Topic 1: An Introduction to Chemistry

Transcription

1 Topic 1: An Introduction to Chemistry Matter & Change (Chapter 1 in Modern Chemistry) Chemistry is a physical science. It is the study of the composition, structure, and properties of matter, the processes that matter undergoes, and the energy changes that accompany these processes. Branches of Chemistry (6 main areas of study) Organic chemistry-the study of most carbon-containing compounds (except carbon & carbonate) Inorganic chemistry-the study of non-organic substances Physical chemistry-the study of the properties and changes of matter and their relation to energy Analytical chemistry-the identification of the components and composition of materials Biochemistry-the study of substances and processes occurring in living things Theoretical chemistry-the use of mathematics and computers to understand the principles behind observed chemical behavior and to design and predict the properties of new compounds A chemical is any substance that has a definite composition. Mass is a measure of the amount of matter. Matter is anything that has mass and takes up space. Matter is stuff. Matter has characteristic properties. These properties can be used to distinguish among substances and to separate them. Properties can be classified as physical or chemical. Physical properties are characteristic that can be observed or measured without changing the identity of the substance. It may look different but it has the same makeup. For example: boiling, melting, and freezing points are physical properties. Water has a formula of H 2 O when it boils and becomes steam which also has a formula of H 2 O. It has this same formula when liquid water freezes and changes to ice. Physical changes are changes in a substance that does not involve a change in the identity of the substance. All phase changes (changes of state) are physical changes. Honors Chemistry Page 1

2 Matter in the solid state has definite volume and definite shape or according to the Kinetic Molecular Theory (KMT), solid particles vibrate around fixed points. Matter in the liquid state has a definite volume but does not have a definite shape. It assumes the shape of its container. According to the (KMT), liquid particles vibrate around moving points. Matter in the gas state has neither definite volume nor definite shape or according to the (KMT), gas particles move in random straight line motion until they hit something. A fourth state of matter is plasma. Plasma is a high-temperature physical state of matter in which atoms lose most of their electrons Physical properties can be intensive or extensive. Extensive properties depend on the amount of matter that is present. For example: volume, mass, and the amount of energy in a substance Intensive properties do not depend on the amount of matter present. For example: melting point, boiling point, density, and ability to conduct electricity and to transfer energy as heat Chemical properties relate to a substance s ability to undergo changes that transform it into different substances. The substances react to form new substances. A chemical change or chemical reaction is a change in which one or more substances are converted into different substances. The substances that react in a chemical change are called the reactants. The substances that are formed by the chemical change are called the products. Task 1a carbon plus oxygen yield carbon dioxide. carbon + oxygen carbon dioxide C + O 2 CO 2 1. Label the following as chemical or physical change/property. a. Melting b. Digestion c. Rusting of iron d. Tearing paper 2. Label the following as an intensive or extensive physical property. a. Length b. Mass c. Density d. Color Honors Chemistry Page 2

3 Classification of Matter Matter is broken down into mixtures and pure substances. Mixtures are blends of two or more kinds of matter, each of which retains its own identity and properties. They are simply physically mixed together and have variable composition. In other words, they do not have a formula. They can usually be separated relatively easily. There are heterogeneous mixtures and homogeneous mixtures. A heterogeneous mixture is not mixed uniformly, for example: granite, clay & water. Homogeneous mixtures are uniform in composition. They still do not have a formula, but the particles are physically mixed evenly all the way through. Homogeneous mixtures are also called solutions. Examples of solutions are salt water, air, and kool-aid. Some mixtures can be separated by filtration, evaporation, distillation, or chromatography. Pure Substances are all homogeneous but are not mixtures because they are chemically combined instead of physically combined. A pure substance has a fixed composition (has a formula). Every sample of a given pure substance has exactly the same characteristic properties and has exactly the same composition. Pure substances can be broken down into elements and compounds. Elements are pure substances that cannot be broken down into simpler, stable substances and is made of one type of atom. Elements are found on the periodic table. They are made of only one type of atom. An atom is the smallest unit of an element that maintains the chemical identity of that element. A compound is a substance that can be broken down into simple stable substances. Each compound is made from the atoms of two or more elements that are chemically bonded. They have definite formulas. Honors Chemistry Page 3

4 Task 1b 1. Label the following as heterogeneous mixture, solution, compound or element. a. Salt, NaCl b. Tea c. Iced tea d. Iron e. Concrete f. Copper g. Sugar The Periodic Table The periodic table is one of the most important references for chemists. It is imperative that you learn the symbols and names of common elements. Be sure to study the list supplied in the memory work section of edline. Periodic table key: blue main group metals purple nonmetals green noble gases pink transition metals teal metalloids Symbol color key (at rt): Black = solid White = gas Yellow = liquid Honors Chemistry Page 4

5 The vertical columns on the periodic table are called groups or families. We refer to the groups by number and the families by name. For example, column 2 is called group 2 or the alkaline earth metal family. Elements in the same group have similar properties. There are a few common family names that you need to know. Other families are known by the first element in that group. Group 1 is the alkali metals family Group 2 is the alkaline earth metal family Group 17 is the halogen family Group 18 is the noble gas family The horizontal rows of elements in the periodic table are called periods. Physical and chemical properties change somewhat regularly across a period. Periods are designated by the number of row (1-7). Elements can also be classified as metals, nonmetals, and metalloids. Metals are to the left of a stair-step line that begins on the left of Boron (except for Hydrogen which is a nonmetal). Nonmetals are to the right of the stair-step line. Elements that touch the stair-step line on a top, bottom or side (not corner) except for aluminum are metalloids. Metals are good conductors of heat and electricity. They are malleable and ductile. They tend to lose valence electrons when they bond. Nonmetals are poor conductors of heat and electricity. They are gaseous or brittle. They tend to gain electrons when they bond. Metalloids sometimes act like metals and sometimes act like nonmetals. (See color codes on periodic table) Task 1c 1. Write the symbol for the following elements. a. Copper b. Sodium c. Argon d. Oxygen e. Zinc 2. Write the name of the following element symbols. a. K b. Li c. S d. He e. Fe Honors Chemistry Page 5

6 3. List the group and period of the following elements. a. Xe b. Cr c. I d. U 4. Label the following as a metal, nonmetal, or metalloid. a. Aluminum b. Nickel c. Germanium d. Carbon Measurements & Calculations (Chapter 2 in Modern Chemistry) The Scientific Method is a logical approach to solving problems by observing and collecting data, formulating hypotheses, testing hypotheses, and formulating theories that are supported by data. Observing uses the senses. Observations can be qualitative or quantitative. Qualitative observations are non-numerical, for example: the room is cold. Quantitative observations are numerical, for example: the room is 58 o F. Chemists study systems. A system is a specific portion of matter in a given region of space that has been selected for study during an experiment or observation. An example of a system is a test tube and its contents. An hypothesis is a testable statement (educated guess). Many times hypotheses are stated as an if-then statement. Hypotheses are tested using an experiment. For an experiment to be valid, the experiment will have controls or conditions that are constant throughout the testing. The experiment will also have variables, or conditions that change. Experiments will usually have a dependent variable and an independent variable. The independent variable (usually on the x axis of a graph) is the variable that is typically being manipulated by the experimenter while the dependent variable (usually on the y axis of a graph) is the observed result of the independent variable being manipulated. An example would be the amount of fertilizer vs. plant growth. The amount of fertilizer would be the independent variable and the growth would be the dependent variable. After the hypothesis has been tested and the data has been recorded and analyzed, then conclusions can be drawn from the experiment. Theorizing about the experiment could start by constructing a model. A model in science is more than a physical object; it is often an explanation of how phenomena occur and how data or events are related (For example: the Honors Chemistry Page 6

7 atomic model). If a model successfully explains many phenomena, it becomes part of a theory. A theory is a broad generalization that explains a body of facts or phenomena. A scientific theory summarizes a hypothesis or group of hypotheses that have been supported with repeated testing. A theory is valid as long as there is no evidence to dispute it. Therefore, theories can be disproven. This is different from a law. A law generalizes a body of observations. At the time it is made, no exceptions have been found to a law. Scientific laws explain things, but they do not describe them. One way to tell a law and a theory apart is to ask if the description gives you a means to explain 'why'. Task 1d 1. Label the following as qualitative or quantitative observations. a. The liquid floats on water. b. The metal is malleable. c. The liquid has a temperature of 55.6 o C. SI Measurements & Units Quantity is something that has magnitude, size, or amount. A measurement is a quantity with a unit of measurement. For example: 1.5 g. Scientists us the SI system of measurement (Le Systeme International d Unites). We will also refer to this as the metric system. Here is a table of the base (fundamental) units used in the SI system. We will only be using the first five. Important fundamental measurements Mass- a measure of the quantity of matter, measured with a balance, units g, kg, etc. Length- the distance between two points, measure with a ruler, units m, km, cm, mm, etc. Honors Chemistry Page 7

8 Derived SI Units Units not listed on the base unit table are derived units. Derived units are combinations of SI base units. Important derived measurements Weight- a measure of the gravitational pull on matter, measured with a scale, unit N (newtons). Volume- the amount of space occupied by an object, measures with a ruler or graduated cylinder, units m 3, cm 3, ml, L, etc. For regular shape objects, you can use a mathematical formula for volume such as V = l x w x h IMPORTANT: 1 cm 3 = 1 ml and 1 dm 3 = 1 L. Density- the ratio of mass to volume, or mass divided by volume. D = m/v, the units for density are g/cm 3, g/ml, kg/l, kg/dm 3, etc. D = m V Task 1e 1. Label each of the following measurements by the quantity each represents. For instance, a measurement of 10.6 kg/m 3 represents density. a. 5.0 g/ml b. 37 s c g d. 47 J e cm 3 f. 500 m 2 g ml Honors Chemistry Page 8

9 Task 1f 1. What is the density of a block of marble that occupies 310. ml and has a mass of 853 g? 2. Diamond has a density of 3.26 g/cm 3. What is the mass of a diamond that has a volujme of cm 3? 3. What is the volume of a sample of liquid mercury that has a mass of 76.2 g, given that the density of mercury is 13.6 g/ml? 4. A block of sodium that has the measurements 3.00 cm x 5.00 cm x 5.00 cm has a mass of 75.5 g. Calculate the density of sodium. Dimensional Analysis (Factor Label Method) Dimensional Analysis is a mathematical technique that allows you to use units to solve problems involving measurements. THIS IS AN IMPORTANT TECHNIQUE! Use conversion factors to go from the quantity given to the quantity sought. Factors are numbers, labels are units. When using the factor-label method, problems consist of three parts: 1. a known beginning GIVEN 2. a desired end WANTED 3. a connecting path CONVERSION FACTORS A conversion factor is a ratio derived from the equality between two different units that can be used to convert from one unit to the other. For example: This can be written as two conversion factors: 4 quarters = 1 dollar 4 quarters or 1 dollar 1 dollar 4 quarters Notice that each conversion factor equals 1. They equal each other. Notice conversion factors can be flipped. Here is an example using conversion factors & dimensional analysis: How many seconds are in 2.5 days? 2.5 days 24 hours 60 minutes 60 seconds 1 day 1 hour 1 minute Multiply the top, multiply the bottom, and divide the answers. All units cancel out except seconds. This equals seconds. Honors Chemistry Page 9

10 Metric Conversions Learn this chart! (Here s an easy way: My kangaroo has dance until dawn cause music makes noise) M k h da u d c m n Mega- kilo- hecta- deka- deci- centi- milli- micro- nano- You can use dimensional analysis with this information or move the decimal the appropriate number of places. For example: Convert 35 m to km. 35 m 1 km = km 1000 m Or m has no prefix so it is in the units column, km has kilo as its prefix, so it is in the k column. Move the decimal 3 places to the left, just like on the chart. If you use this method, watch out for prefixes that change by more than one change of 10 (mega, micro, nano). 35 m becomes km Task 1g 1. Complete the following conversions a g = kg b km = m c g = g d. 3.5 mol = mol e. 1.2 L = ml f ml = cm 3 2. Use dimensional analysis to determine the following. a. If 1 mag = 13 bops and 1 bop = 4.6 skuts, how many mags are in 583 skuts? b. How many centimeters are in 2.5 yards? (1 inch = 2.54 cm) Honors Chemistry Page 10

11 Temperature Conversions Use the following formulas for temperature conversions. Task 1h K = o C o C = K 273 o C = 5/9( o F-32) o F = 9/5 o C Make the following temperature conversions. a o C = K b. 323 K = o C c. 75 o F = o C d. 368 K = o F Honors Chemistry Page 11

12 Using Scientific Measurements Accuracy refers to the closeness of measurements to the correct or accepted value of the quantity measured. Precision refers to the closeness of a set of measurements of the same quantity made in the same way. Percentage Error is calculated by subtracting the accepted value from the experimental value, dividing the difference by the accepted value, and then multiplying by 100. Percentage error = experimental - accepted accepted x 100 Task 1i 1. What is the percentage error for a mass measurement of 17.7 g, given that the correct value is 21.1 g? 2. A volume is measured experimentally as 4.26 ml. What is the percent error, given that the correct value is 4.15 ml? 3. During an experiment, a student obtains the following density data: 9.12 g/ml, 9.11 g/ml, and 9.13 g/ml. The literature value for the density of this substance is 8.78 g/ml. Is this student accurate? Is this student precise? What is the % error of this students information? Honors Chemistry Page 12

13 Significant figures (Significant Digits) in a measurement consist of all the digits known with certainty plus one final digit, which is somewhat uncertain or is estimated. Rules for Significant Figures Read from the left and start counting sig figs when you encounter the first non-zero digit 1. All non zero numbers are significant (meaning they count as sig figs) 613 has three sig figs has six sig figs 2. Zeros located between non-zero digits are significant (they count) 5004 has four sig figs 602 has three sig figs has 16 sig figs! 3. Trailing zeros (those at the end) are significant only if the number contains a decimal point; otherwise they are insignificant (they don t count) has four sig figs has six sig figs has two sig figs unless you re given additional information in the problem 4. Zeros to left of the first nonzero digit are insignificant (they don t count); they are only placeholders! has three sig figs has two sig figs also has two sig figs! Rules for addition/subtraction problems Your calculated value cannot be more precise than the least precise quantity used in the calculation. The least precise quantity has the fewest digits to the right of the decimal point. Your calculated value will have the same number of digits to the right of the decimal point as that of the least precise quantity. In practice, find the quantity with the fewest digits to the right of the decimal point. In the example below, this would be 11.1 (this is the least precise quantity) = (this is what your calculator spits out) In this case, your final answer is limited to one sig fig to the right of the decimal or 25.3 (rounded up). Honors Chemistry Page 13

14 Rules for multiplication/division problems The number of sig figs in the final calculated value will be the same as that of the quantity with the fewest number of sig figs used in the calculation. In practice, find the quantity with the fewest number of sig figs. In the example below, the quantity with the fewest number of sig figs is 27.2 (three sig figs). Your final answer is therefore limited to three sig figs. (27.2 x 15.63) = (this is what you calculator spits out) In this case, since your final answer it limited to three sig figs, the answer is 230. (rounded down) Rules for combined addition/subtraction and multiplication/division problems First apply the rules for addition/subtraction (determine the number of sig figs for that step), then apply the rules for multiplication/division. Task 1j 1. Provide the number of sig figs in each of the following numbers: (a) g (b) 3.40 x 103 ml (c) g (d) L (e) g (f) 1020 L 2. Perform the operation and report the answer with the correct number of sig figs. (a) (10.3 m) x ( m) = (b) (10.3) + ( ) = (c) [(10.3) + ( )] = [(10.3) x ( )] 3. Polycarbonate plastic has a density of 1.2 g/cm 3. A photo frame is constructed from two 3.0 mm sheets of polycarbonate. Each sheet measures 28 cm by 22 cm. What is the mass of the photo frame? Honors Chemistry Page 14

15 Scientific Notation (You should already know this) Scientific Notation was developed in order to easily represent numbers that are either very large or very small. Scientific Notation is based on powers of the base number 10. Examples: The number 200,000,000,000 stars in scientific notation is written as 2 x stars The number ,006,645 kilograms in scientific notation is written as x 10-6 stars The first number is called the coefficient. The coefficient must be greater than or equal to 1 and less than 10. The coefficient contains only significant digits. The second number is called the base. The base must always be 10 in scientific notation. The number -6 is referred to as the exponent or power of ten. The exponent must show the number of places that the decimal needs to be moved to change the number to standard notation. A negative exponent means that the number written in standard notation is less than one. To Change from Standard Form to Scientific Notation: 1. Place decimal point such that there is one non-zero digit to the left of the decimal point. 2. Count number of decimal places the decimal has "moved" from the original number. This will be the exponent of the If the original number was less than 1, the exponent is negative; if the original number was greater than 1, the exponent is positive. To Change from Scientific Notation to Standard Form: 1. Determine the number of places the decimal must be moved from the exponent. 2. Decide if the standard form will be a number greater than one or less than one. 3. Move the decimal in the coefficient adding place holders if necessary. Task 1k 1. Write the following numbers in scientific notation. a b c d Honors Chemistry Page 15

16 2. Write the following numbers in standard notation. a x 10 3 b x 10-1 c x 10 9 d. 512 x 10-8 Data Manipulation All measurements taken in lab must be in the correct significant digits. Data tables should be prepared ahead of time if possible. You will be expected to understand what is occurring, not only following directions. This is not just cookbook chemistry. You will have to graph data that you collect in lab. Remember that the independent variable goes on the x-axis. The dependent variable goes on the y-axis. Two quantities are directly proportional to each other if dividing one by the other gives a constant value. y/x = k or y = kx (in the form of y = mx + b) This is a straight line graph. An example of this is density. Two quantities are inversely proportional to each other if their product is constant. xy = k This graph produces a curve called a hyperbola. Pressure & volume of gases give this type of inverse proportion. States of Matter (Chapter 10 Modern Chemistry) Watch this!! This will review the six phase changes that matter undergoes. Phase Diagrams A phase diagram is a graph of pressure versus temperature that shows the conditions under which the phases of a substance exist. The triple point of a substance indicates the temperature and pressure conditions at which the solid, liquid, and vapor of the substance can coexist at equilibrium. The critical point of a substance indicates the critical temperature and critical pressure. This is the point above which the substance cannot exist in the liquid state. Honors Chemistry Page 16

17 The normal melting point or boiling point is the melting or boiling point at standard pressure. In order to find the normal points you must read the graph across from 1 atm, 760 mmhg, or kpa. General Phase Diagram Honors Chemistry Page 17

18 Task 1L 1. Using phase Diagram for water (a), what is the critical point of water? 2. What is water s triple point? 3. Using the phase diagram for carbon dioxide (b), what is the normal melting point? The normal boiling point? Heating & Cooling Curves Heating & Cooling curves can help you tell at what temperature a substance melts/freezes or boils/condenses at the current pressure. Note that as a phase is change the temperature doesn t change. All the energy is going toward breaking the particles intermolecular bonds so the change can occur. After the phase change occurs the energy can now be used to increase the temperature (or vice versa). This link allows you to perform three melting/boiling experiments while simultaneously graphing the data. You should be able to determine the melting point and boiling point of each substance. Honors Chemistry Page 18

19 Solutions (Chapter 12 Modern Chemistry) Solubility & Solubility Curves Solubility of a substance is the amount of that substance required to form a saturated solution with a specific amount of solvent at a specified temperature. A saturated solution contains the maximum amount of dissolved solute. An unsaturated solution contains less solute than a saturated solution under the existing conditions. A supersaturated solution contains more dissolved solute than a saturated solution contains under the same conditions. As you can see in the diagram below, the solubility of solids generally increase with temperature, while the solubility of gases decrease with a temperature increase. Task 1m 4. Using the solubility curves, what is the solubility of KClO 3 in 100 g of H 2 O at 50 o C? 5. How many grams of KCl is needed to make a saturated solution at 50 o C in 100 g of H 2 O? What about in 200 g of H 2 O at this same temperature? 6. How many grams of NaCl will dissolve in 100 g of H 2 O at 60 o C? How much salt would sink to the bottom of the beaker if 30 g of NaCl is added to 100 g of H 2 O? What is 60 g of NaCl is added to 100 g of H 2 O? 7. What type of solution would 40 g of NH 3 in 100 g H 2 O at 20 o C? Honors Chemistry Page 19

20 Topic 2: Atomic Theory & Structure Atoms: The Building Blocks of Matter (Chapter 3 in Modern Chemistry) History of Atom Theory Democritus a great Greek thinker, ( BC) first proposed the existence of an ultimate particle. He used the word "atomos" to describe this particle. He also explain solids by stating that their atoms must be like sandspurs, so they stuck together while liquid atoms were more like marbles that could roll over each other. Democritus Aristotle ( BC) was a proponent of the continuum. He believed in the four elements of air, earth, water and fire. Aristotle felt that regardless of the number of times you cut a form of matter in half, you would always have a smaller piece of that matter. This view held sway for 2000 years primarily because Aristotle was the tutor of Alexander the Great. In other words, he was more popular than Democritus. Antoine Lavoisier ( ) was the first person to make good use of the balance. He was an excellent experimenter and he based his ideas on experimental evidence. He proposed the Law of Conversation of Mass which represents the beginning of modern chemistry. This Law of Conservation of Mass states that matter cannot be created or destroyed during normal chemical reactions or physical changes, it merely rearranges or changes form. Soon after Lavoisier s discovery, Joseph Proust ( ) developed the law of definite proportions. This law stated that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. In other words, a compound like salt, NaCl, is made of the same ratio of Na to Cl by mass, whether the salt is mined from the ground, collected from the sea, or produced in the lab. John Dalton ( ) proposed the Law of Multiple Proportions. This law led directly to the proposal of the Atomic Theory in The Law of Multiple Proportions states that if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers. Honors Chemistry Page 1

21 Dalton's Atomic Theory 1) All matter is made of extremely small particles called atoms. Atoms are indivisible and indestructible. 2) All atoms of a given element are identical in mass and properties. Atoms of different elements differ in mass and properties. 3) Compounds are formed by a combination of two or more different kinds of atoms in simplewhole number ratios. 4) A chemical reaction is a rearrangement of atoms. Modern atomic theory is, of course, a little more involved than Dalton's theory but the essence of Dalton's theory remains valid. Today we know that atoms can be destroyed via nuclear reactions but not by chemical reactions. Also, there are different kinds of atoms (differing by their masses) within an element that are known as "isotopes", but isotopes of an element have the same chemical properties. Many heretofore unexplained chemical phenomena were quickly explained by Dalton with his theory. Dalton's theory quickly became the theoretical foundation in chemistry. History of Atomic Structure John Dalton thought that atoms were the smallest possible particle. We now know that the atom is made up of even smaller particles. Here are the experiments that led to the discovery of what we believe is the structure of the atom. J. J. Thomson ( ) identified the negatively charged electron in the cathode ray tube in He deduced that the electron was a component of all matter and calculated the charge to mass ratio for the electron. He also proposed the "plum pudding" model of the atom. In this model, the volume of the atom is composed primarily of the more massive (thus larger) positive portion (the plum pudding). The smaller electrons (actually, raisins in the plum pudding) are dispersed throughout the positive mass to maintain charge neutrality. Honors Chemistry Page 2

22 Robert Millikan ( ) determined the unit charge of the electron in 1909 with his Oil Drop experiment at the University of Chicago. This allowed for the calculation of the mass of the electron. In this experiment, he used charged plates to balance droplets of oil to determine the charge to mass ratio of the droplets. Ernest Rutherford ( ) proposed the nuclear atom as the result of the gold-foil experiment in Rutherford's Gold Foil Experiment What happened What it means Alpha particles (positive helium nuclei) were shot at a piece of gold foil. Most of the particles went straight through. A few particles were slightly deflected A very few particles were reflected, bounced back. Most of the atom is empty space. Since alpha particles are +, they must have come close to + protons. The alpha particles hit something dense. This is the discovery of the nucleus. Honors Chemistry Page 3

23 Rutherford proposed that all of the positive charge and all of the mass of the atom occupied a small volume at the center of the atom called the nucleus and that most of the volume of the atom was empty space occupied by the electrons. Francis Aston ( ) invented the mass spectrograph in He was the first person to observe isotopes. For example he observed that there were three different kinds of hydrogen atoms. James Chadwick ( ) discovered the neutron in Chadwick was a collaborator of Rutherford's. Interestingly, the discovery of the neutron led directly to the discovery of fission and ultimately to the atomic bomb. Neils Bohr ( ) contributed ideas to the current atomic structure models. Bohr stated that the electrons were in energy levels in which the electrons were quantized. In other words, electrons on different energy levels have different amounts of energy. The current atomic model is the electron cloud/probability model. The energy levels are 3 dimensional volumes of electrons around the nucleus instead of rings. Honors Chemistry Page 4

24 Atomic Structure Summary The atom is made of a nucleus that contains protons (+) and neutrons (neutral). Surrounding the nucleus are energy levels that hold electrons (-). The atom itself is neutral because there are equal number of protons and electrons. The nucleus is positive. Properties of Subatomic Particles Particles Symbols Relative charge Mass Number Location Electron e -, 0-1e -1 0 amu Energy levels Proton p +, 1 1H 1 1 amu Nucleus Neutron n 0, 1 0n 0 1 amu Nucleus Note: The unit amu stands for atomic mass unit. It equals 1/12 of the carbon-12 atom. Task 2a 1. List the scientist that developed or discovered the following. a. Discovered the nucleus b. Discovered the electron c. Oil Drop experiment d. Discovered the neutron Counting Atoms Atoms of different elements have different numbers of protons. Atoms of the same element all have the same number of protons. The atomic number (Z) of an element is the number of protons of each atom of that element. The atomic number is found on the periodic table. In this particular square of gold from the periodic table the top number is the atomic number. It is always a whole number. The periodic table is arranged by increasing atomic number. This means that gold has an atomic number of 79 and has 79 protons. The atomic number identifies the element. Any particle that has 79 protons is gold no matter how many electrons or neutrons it has. Remember atoms have a neutral charge so if the gold atom has 79 protons, it also has 79 electrons. The bottom number of this particular square from a periodic table is the atomic mass or mass number. The mass number (A) is the total number of protons and neutrons that make up the nucleus of an atom. The mass number does not have to be same for every atom of an element. Atoms of the same element that have different masses are called isotopes. Since electrons affect chemical properties, different isotopes of the same element have the same chemical properties. The symbols for isotopes can be written as hyphen notation (gold-197) or nuclear notation Honors Chemistry Page 5

25 ( Au). In hyphen notation the mass number is given. You have to look up the atomic number on the periodic table. In nuclear notation, the top number is the mass number and the bottom number is the atomic number. Notice that mass numbers on the periodic table are not whole numbers. They are weighted averages of all the isotopes of that particular element. You can also calculate the protons, electrons and neutrons for ions. Ions are atoms that have gained or lost electrons. They have a charge. Charges are written as a superscript on the right of the symbol Na 1+ Atomic particle summary Atomic number = the number of protons in the nucleus of one atom of the element Mass number = the number of protons and neutrons Example: # of protons = the atomic number # of electrons = the # of protons for atoms, if an ion add electrons for a charge and subtract electrons for a + charge # of neutrons is the mass number atomic number Calculate the number of protons, electrons, and neutrons in the following P protons = 15 electrons = 15 neutrons = 16 Na-23 protons = 11 electrons = 11 neutrons = Se 2- protons = 34 electrons = 36 neutrons = Ag + protons = 47 electrons = 46 neutrons = 61 Honors Chemistry Page 6

26 Task 2b 1. Consider the following pairs; does either pair represent a pair of isotopes? a K and 40 18Ar b Sr and 94 38Sr 2. Determine the number of protons, electrons, and neutrons in the following. a Pb b S c. Br-80 d Cu 2+ e. 15 7N 3-3. Write the hyphen notation and the nuclear notation for an isotope with 15 electrons and 15 neutrons. 4. Write the nuclear notation for an ion that has 53 protons, 54 electrons, and 73 neutrons. Average Atomic Mass Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element. Average atomic mass = Sum of (% of each isotope)(atomic mass of each isotope) 100 Task 2c 1. Three isotopes of argon occur in nature 36 18Ar, 38 18Ar, and 40 18Ar. Calculate the average atomic mass of argon to two decimal places, given the following relative atomic mass and abundances of each of the isotopes: argon-36 (35.97 amu; 0.337%), argon-38 (37.96 amu; 0.063%), and argon-40 (39.96 amu; %). 2. Naturally occurring boron is 80.20% boron-11 (atomic mass amu) and 19.80% of some other isotopic form of boron. What must the atomic mass of this second isotope be in order to account for the amu average atomic mass of boron? (Write the answer to two decimal places.) Honors Chemistry Page 7

27 Topic 3: Electromagnetic Spectrum & Quantum Theory Arrangement of Electrons in Atoms (Chapter 4 in Modern Chemistry) Introduction There were two light theories in the early Sir Isaac Newton subscribed to the particle theory of light. Christian Huygens subscribe to the wave theory of light. There was data to support both theories. Einstein developed the idea of a photon, Bohr proposed a quantized model of the atom, and eventually, Louis debroglie, came up with the Wave-Particle Duality of Nature. He said that sometimes waves act like particles and sometimes particles act like waves. This was true of very small particles. Eventually, the Quantum theory was developed and so were electron configurations. Newton Huygens Einstein Bohr debroglie Properties of Light Visible light is a kind of electromagnetic radiation, which exhibits wavelike behavior as it travels through space. There are other types of light that make up the electromagnetic spectrum. All forms of electromagnetic radiation move at a constant speed. We call this the speed of light (c); c = 3.00 x 10 8 m/s. Honors Chemistry Page 1

28 Wave motion is repetitive. It is characterized by wavelength and frequency. Wavelength ( ) is the distance between corresponding points on adjacent waves. Wavelength is a length unit so it is expressed in meters, centimeters, or nanometers. Frequency ( is defined as the number of waves that pass a given point in a specific time, usually one second. Frequency is expressed in waves/second, s -1,or hertz (Hz). Wavelength & frequency are inversely proportional, meaning longer wavelengths have lower frequencies and vice versa. Mathematically, they are related to each other in the following relationship. c = The Photoelectric Effect The photoelectric effect refers to the emission of electrons from a metal when light shines on the metal. It was found that only light above certain minimum frequencies could cause these electrons to be released no matter how intense the light. This could not be explained by the wave theory. Max Planck suggested that energy is emitted or absorbed not as continuous waves, but as small, specific packets of energy called quanta. A quantum of energy is the minimum quantity of energy that can be lost or gained by an atom. E = h E is the energy, in joules is the frequency in s -1 or Hz h is Planck s constant;; h = 6.63 x Js or J/Hz Max Planck By combining the equations, c = and E = h, another useful equation can be derived. E = hc Honors Chemistry Page 2

29 Task 3a 1. Characterize each of the following as absorption or emission: a. an electron moves from E 2 to E 1 b. an electron moves from E 1 to E 3 c. an electron moves from E 6 to E 3 2. Which energy-level change above emits or absorbs the highest energy? The lowest energy? 3. Solve the following problems using the equations in this section. a. Determine the frequency of light whose wavelength is x 10-7 cm. Rydberg Equation b. Determining the energy in joules of a photon whose frequency is 3.55 x Hz. c. When sodium is heated; a yellow spectral line whose energy is 3.37 x J per photon is produced. What is the wavelength of this light? d. The laser in an audio compact disc player uses light with a wavelength of 7.80 x 10 2 nm. Calculate the frequency of this light. Calculate the energy of a single photon of this light. You can also determine the amount of energy absorbed or emitted as electrons move from one energy level to another using a form of the Rydberg equation. E = E f E i = E photon E = x n 2 E = ( x J ) ( 1 n f 2-1 n i 2 ) E = the energy associated with a particular quantum number, n. By calculating the energies associated with two different quantum levels and finding the difference, one can calculate the energy required to promote an electron from one level to another, or calculate the energy released when an electron falls back from a higher level to a lower level. Energy emitted is exothermic and is given a negative charge. Energy absorbed is endothermic and is given a positive charge. Task 3b 1. What is the energy difference when an electron moves from E 3 to E 5? Is the energy emitted or absorbed. Use the appropriate sign. Honors Chemistry Page 3

30 The Hydrogen-Atom Line-Emission Spectrum When electricity is passes through a gas at low pressure, the energy of some of the gas atoms increases. The lowest energy state of an atom is its ground state. A state in which an atom has a higher energy than it has in its ground state is an excited state. For an electron to move from its ground state to an excited state, energy must be absorbed. When an electron falls from an excited state to a lower excited state, or to its ground state, energy will be released. This energy is released in the form of light called a photon. The production of colored light in neon lights is an example of this process. When electricity was passed through hydrogen gas at low pressure, a pinkish glow was given off. When a narrow beam of the emitted light was shined through a prism, it was separated into four specific colors of the visible spectrum. The four bands (lines) of light were part of what is known as hydrogen s line-emission spectrum. These distinct bands at specific wavelength & frequencies indicated that the energy differences between the atoms energy states were fixed. This led to the Bohr model of the atom. In this model, Niels Bohr linked the atom s electron to photon emission. The electron can circle the nucleus only in allowed paths, or orbits. When the electron is in one of these orbits, the atom has a definite, fixed energy. Here are two animations that will help you visualize what happens to electrons as they absorb or emit energy. The Quantum Model of the Atom So, after the photoelectric effect and hydrogen s line-emission spectrum revealed that light could behave as particles and waves, Louis de Broglie proposed the wave-particle duality of nature. Then Werner Heisenberg tried to determine where the electrons were in the atom. Because you need a photon to Honors Chemistry Page 4

31 detect an electron, and a photon causes an electron to be knocked off course, there is always an uncertainty about the location of an electron. This is called the Heisenberg uncertainly principle. It states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle. There are only probable positions of electrons. Max Planck founded the Quantum theory, which describes mathematically the wave properties of electron and other very small particles. This theory states that electrons do not travel around the nucleus in neat orbits, as Bohr postulated, but exist in certain regions or volumes, called orbitals. An orbital is a threedimensional region around the nucleus that indicates the probable location of an electron. Atomic Orbitals and Quantum Numbers s orbital In order to completely describe orbitals, scientists use quantum numbers. Quantum numbers specify the properties of atomic orbitals and the properties of electron in those orbitals. There are four quantum numbers. The first three indicate the main energy level, the shape, and the orientation of an orbital. The fourth, describes the spin on an electron in the orbital. Quantum number notes There are 4 Quantum Numbers: n,, m, s 1. n = Principal Quantum Number (n = 1, 2, 3 ) Represents the main energy level, if n = 1, 1 st energy level, n = 2, 2 nd energy level, etc. This also represents the size of the electron cloud. n = 1 is the energy level closest to the nucleus, the larger n, the farther away from the nucleus 2n 2 determines maximum number of electrons that can occupy an energy level 2. = sublevel ( = 0 to n-1) Refers to different energy states in each energy level Determines the shape of the orbital(s) Number of sublevels = n o n = 1, has 1 sublevel s ( = 0) o n = 2, has 2 sublevels s ( = 0), p ( = 1) Honors Chemistry Page 5

32 o n = 3, has 3 sublevels s ( = 0), p ( = 1), d ( = 2) o n = 4, has 4 sublevels s ( = 0), p ( = 1), d ( = 2), f ( = 3) 3. m = orbital (m = - to + ) Space occupied by a pair of electrons Determines orientation in space Each orbital can be represented by a box Each orbital can hold a maximum of two electrons s = 1 orbital, 1 box, a maximum of 2 electrons (0) p = 3 orbitals, 3 boxes, a maximum of 6 electrons (-1,0,1) d = 5 orbitals, 5 boxes, a maximum of 10 electrons (-2, -1, 0, 1, 2) f = 7 orbitals, 7 boxes, a maximum of 14 electrons (-3, -2, -1, 0, 1, 2, 3) Honors Chemistry Page 6

33 4. s = spin in order for 2 electrons to occupy the same orbital, they must have opposite spins the box represents the orbital, the arrows represent electrons spinning in opposite directions, there cannot be arrows pointing in the same direction in the same box Each arrow represents an electron (+ ½, -½) Task 3c Answer the following questions based on quantum numbers. 1. How many quantum numbers are there? 2. What is the maximum number of electrons in an orbital? 3. What is the maximum number of electrons in n = 2? 4. Which quantum number represents the shape of the electron cloud? 5. Which quantum number represents the volume of the electron cloud? 6. How many orbitals are in the p sublevel? 7. How many electrons can be in the d sublevel? 8. Which energy level has the lowest energy? 9. Which sublevel has the lowest energy? 10. How many sublevels are in n = 4? What are they? Rules for placing electrons in orbitals Pauli s Exclusion Principle No two electrons in an element can have the same set of quantum numbers. Aufbau Principle Electrons occupy orbitals of lowest energy first 1s Hund s Rule Within a sublevel, orbitals are half-filled with electrons before they become filled p sublevel with 3 electrons Honors Chemistry Page 7

34 n m s Honors (ex.) 1 s 1 1,0,0,+ ½ 1,0,0,- ½ 2 s, p 1, 3 2,0,0,+ ½ 2,0,0,- ½ 3 s, p, d 1, 3, 5 2,1,-1, + ½ 2,1,-1, - ½ 2,1,0, + ½ 2,1,0, - ½ 2,1,1, + ½ 2,1,1, - ½ 4 s, p, d, f 1, 3, 5, 7 Look up the atomic number of the element for which you are drawing the orbital filling diagram. Since these are atoms, this not only tells you the number of protons but also the number of electrons. The electrons are represented by arrows. For example, K has atomic number of 19, so its atoms have 19 electrons, therefore, 19 arrows. For example the orbital filling diagrams for for H is: For He: 1s 1 1s 2 For Li: For Be: 1s 2 2s 1 1s 2 2s 2 Honors Chemistry Page 8

35 For B: For N: 1s 2 2s 2 2p 1 1s 2 2s 2 2p 3 For F: 1s 2 2s 2 2p 5 For Si: 1s 2 2s 2 2p 6 3s 2 3p 2 Note that you always start with 1s. Always start with an up arrow. When you are starting to fill up the orbitals (boxes) you placed up arrows in all the boxes until there is at least one arrow in the orbital, then you go back and finish filling that orbital. Only two arrows can go in a box. You can also write the electron configurations for the same elements. You just write the numbers below without drawing the boxes. For H: 1s 1 (read as: one s one) For He: 1s 2 For Li: 1s 2 2s 2 For Be: 1s 2 2s 2 For B: 1s 2 2s 2 2p 1 For N: 1s 2 2s 2 2p 3 For F: 1s 2 2s 2 2p 5 For Si: 1s 2 2s 2 2p 6 3s 2 3p 2 It is very important that you write the electron configuration in order of lowest energy to highest energy. You can use the diagonal rule or the periodic table. The diagonal rule shows the order of electron filling. Follow the arrows. Honors Chemistry Page 9

36 Task 3d 1. Draw the orbital filling diagram for the following. a. P b. V c. Na 2. Write the electron configuration for the following. a. W b. Zn c. Ca d. Tl e. Br 3. What are the quantum numbers for the last electron in each of the following? a. As b. Nb c. Ba 4. Four electrons in an atom have the four sets of quantum numbers given below. Which electrons are in the same orbital? Explain your answer. a. 1, 0,0, - b. 1, 0, 0, + c. 2, 1, 1, + d. 2, 1, 0, + 5. Which of the sets of quantum; numbers below are possible? Which are impossible? Explain your choices. a. 2, 2, 1, + b. 2, 0, 0, - c. 2, 0, 1, - You also need to be able to write the quantum numbers for electrons. In the notes above there were numbers in parentheses. These are the quantum numbers. For example: Li has an electron configuration of the quantum numbers for the last electron would be 2, 0, 0, + ½. 1s 2 2s 1 Honors Chemistry Page 10

37 The 2 represents the 2 level, the first 0 represents the s sublevel, the second 0 represents the box, and the + ½ represents the clockwise spin. (We actually do not know if it is clockwise or counterclockwise, by convention we use + ½ for the up arrow.) Another example: Nitrogen s orbital filling diagram is 1s 2 2s 2 2p 3 The quantum numbers for the last electron would be 2, 1, -1, + ½. Electron Configuration & the Periodic Table (Chapter 5 in Modern Chemistry) You can also write electron configurations using the periodic table. Below is a periodic table labeled with the appropriate blocks. This is the method I use, and the method taught in most college classes The rows on the periodic table represent the outer energy level or principle quantum number. The first two columns are the s blocks plus He. The middle section (transition metals) is the d block elements. Notice that the principle quantum number decreases by one in this section. In other words a d block element on row 5, has a 4d sublevel. The last 6 columns of the periodic table are the p block elements. Notice that the principle quantum number for these elements is the same as the row number. The f section of the periodic table is located on the bottom two rows of the table. Honors Chemistry Page 11

38 When using the periodic table to write electron configurations, follow the chart, counting the elements in each block until you reach the element you are writing the configuration for. Alternately, you can write the noble gas configuration, by writing the noble gas immediately before the specified element in brackets [ ]. Here are some examples using the periodic table above. Electron Configuration Noble gas configuration For Sr: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 [Kr] 5s 2 For Al: 1s 2 2s 2 2p 6 3s 2 3p 1 [Ne] 3s 2 3p 1 For Pb: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 2 [Xe] 6s 2 4f 14 5d 10 6p 2 Now you try some. Task 3e 1. Write the quantum numbers for the last filling electron in the following. a. K b. Co c. Sn 2. Write the electron configuration for the following using the periodic table. a. O b. Mo c. Cs 3. Write the noble gas configuration for the following. a. Mg b. Fe c. Cl d. Au e. Fr Honors Chemistry Page 12

39 4. Write the symbol of the element that is represented below by the electron configuration or noble gas configuration. a. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 b. [Kr] 5s 2 4d 10 5p 3 c. 1s 2 2s 2 2p 6 3s 2 3p 4 Exceptions to the electron configuration There are some elements that do not follow the general rules for writing electron configuration, for example, Cr and Cu. The anomalies of Cr and Cu are easy to explain once you know that a half-filled or completely filled d shell is considered to have extra stability. Hence configurations ending 4s 1 3d 5 and 4s 1 3d 10 rather than 4s 2 3d 4 and 4s 2 3d 9 are considered to be preferable. In each case one of the s electrons is promoted to the d shell to create a more stable configuration. Valence Electrons Valence electrons are the electrons that are located on the outside energy level. You can tell which electrons are valence by the main energy level. For example, in the electron configuration, 1s 2 2s 2 2p 6 3s 2 3p 4, the highest energy level is 3. There are 6 electrons in level three, two in the s sublevel, and 4 in the p sublevel. Electron Dot Diagrams Electron Dot Diagrams are visual representations of the valence electrons in an atom. For example: Al has a noble gas configuration of [Ne] 3s 2 3p 1. If I were to draw the orbitals for the valence electrons only they would be. 3s 3p There are two electrons in the 3s and 1 electron in the first orbital of the 3p. I could draw the valence electrons around the symbol. Al The electron dot diagram for Al would be : Al. It doesn t matter where you put the electrons only that you put the correct number of electrons in each box. It is called electron dot because you use dots instead of arrows. In the electron dot diagram, the symbol represent the nucleus and all the inner electrons. The dots represent the valence electrons. Honors Chemistry Page 13

40 Here s another example: Te [Kr] 5s 2 4d 10 5p 4 (Remember 4d is not valence) 5s 5p The electron dot would be Task 3f 1. Using the periodic table, how many valence electrons are in the following elements? a. Br b. Sr c. Ar d. Fr 2. Draw the electron dot diagram for each atom in (1). Electron Configurations of Ions You can also write the electron configurations of ions. The positive ions have lost electrons. Electrons can only be removed from the outside energy level. For example, Na has an electron configuration of [Ne] 3s 1. Na + has an electron configuration of [Ne] or preferably [He] 2s 2 2p 6. Fe has an electron configuration of [Ar] 4s 2 3d 6. Fe 2+ has an electron configuration of [Ar] 3d 6. Fe 3+ has an electron configuration of [Ar] 3d 5. Negative ions have gained electrons. Electrons should be placed in the next available orbital. For example: S has and electron configuration of [Ne] 3s 2 3p 5. Cl - has an electron configuration of [Ne] 3s 2 3p 6. Isoelectronic Species Species that have the same electron configuration are called isoelectronic. These usually come in a series. For example: Ne is isoelectronic with O 2-, F -, Na +, and Mg 2+. Predicting Oxidation Numbers An oxidation number is the tendency of an atom to gain or lose electrons. According to the octet rule atoms are more stable whenever they have 8 valence electrons (except for things in period 1). To determine the oxidation number, you need to know how many electrons are gained or lost, or how many are likely to be gained or lost. Honors Chemistry Page 14

41 For example: Ca has a configuration of [Ar] 4s 2. It is easier to lose those 2 valence electrons than to gain 6 to make 8. Losing electrons cause the charge to be positive, so the oxidation number of Ca that is most likely is 2 +. The calcium ion is usually written as Ca 2+. P has a configuration of [Ne] 3s 2 3p 3. It has 5 valence electrons. It is easier to gain three electrons to make a total of eight than to lose the 5 it has. So it will gain 3 electrons, making it a negative three charge. This is written as P 3-. Task 3g 1. Without looking at the periodic table, identify the group, period, and block in which the element that has the electron configuration [Xe] 6s 2 is located. 2. Without looking at the periodic table, write the electron configuration for the Group 1 element in the third period. Is this element likely to be more reactive or less reactive than the element described in (1)? 3. Write the electron configurations for the following ions. a. Cl - b. K + c. Fe 2+ d. Fe 3+ e. P 3-4. Which of the ions in (3) are isoelectronic with each other? 5. Which of the following does not have the same configuration as a noble gas: Na +, Rb +, O 2-, Br -, Ca +, Al 3+, S 2-? 6. What is the probable oxidation number for the following elements? a. Mg b. K c. S d. I e. P 7. What is the electron configuration of silver? It only has one oxidation number, what is it? Honors Chemistry Page 15

42 Topic 4: Periodic Table & Trends The Periodic Law (Chapter 5 in Modern Chemistry) The History of the Periodic Table Stanislao Cannizzaro discovered a method of accurately measuring atomic masses. Dmitri Mendeleev is credited with organizing the first periodic table based on atomic masses.(1869) He noticed the when the elements were arranged in order of increasing atomic mass, certain similarities in their chemical properties appeared at regular intervals. These repeating patterns are referred to as periodic. Mendeleev did find some discrepancies, he placed iodine after tellurium even though based on atomic masses they should be reversed. Tellurium acted more like O, S, and Se. Iodine acted more like F, Cl, and Br. He knew there was some problems with using atomic masses for ordering the table. He also left blanks in his periodic table. He boldly predicted the existence and properties of the elements that would fill three of these blanks based on the properties of elements that were similar in his table. Eventually, all three of these elements were discovered, Sc, Ga, and Ge. Their properties were strikingly similar to those predicted by Mendeleev. Henri Mosely discovered that the elements in the periodic table fit into patterns better when they were arranged in increasing order according to nuclear charge, or number of protons (atomic number). (1911) This corrected the discrepancies in Mendeleev s table. Periodic Law state the physical and chemical properties of the elements are periodic functins of their atomic numbers. The Modern Periodic Table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group. Remember that the numbers across the top represent the group, while the numbers down the side represent the period. HN Chemistry Page 1

43 The Element Song by Tom Lehrer Noble gases (Group 18) All atoms of the noble gases have their outer s and p orbitals filled. We will see later that these atoms require very large amounts of energy to form ions, so much in fact, that they are difficult to alter chemically and as such are inert (unreactive) and do not tend form ions. Alkali metals (Group 1) & Alkaline Earth metals (Group 2) METALS Group 1 atoms have an electronic structure [Noble gas] ns1. This means that they tend to lose the s electron when they from an ion, leaving behind an inert noble gas type structure. This explains why Group 1 elements tend to only form 1 + ions. Group 2 atoms have an electronic structure [Noble gas] ns2. This means that they tend to lose the two s electrons when they from an ion, leaving behind an inert noble gas type structure. This explains why Group 2 elements tend to only form 2 + ions. A similar argument can be applied to group 3 atoms and their simple ions. HN Chemistry Page 2

44 Groups 16 & 17 (Chalcogens & Halogens) NON-METALS Group 16 atoms have an electronic structure [Noble gas] ns2 np4. This means that they tend to gain two p electrons when they from an ion, to reach an inert noble gas type structure with a charge of 2 -. Check out this animation: It is about oxygen. Very cute! Group 17 atoms have an electronic structure [Noble gas] ns2 np5. This means that they tend to gain one p electron when they from an ion, to reach an inert noble gas type structure with a charge of 1-. Task 4a 1. What name is given to each of the following groups of elements in the periodic table? a. Group 1 b. Group 2 c. Groups 3-12 d. Group 17 e. Group Based on what you know about their electron configurations, which groups do you think are the most active? Why? The least active? Why? Periodic Properties As you have learned earlier, the elements in the same group have similar electron configurations, therefore they also have similar physical and chemical properties because their valence electrons are the same. You need to be able to identify the properties and how they relate to each other based on their placement on the periodic table. Before we discuss the periodic property trends, we need to discuss two reasons that properties change based on the periodic table. The shielding effect causes properties to change that are in the same group (column). The shielding effect occurs because the inner electrons shield the nucleus from the outer electrons. The more inner electrons there are between the valence electrons and the nucleus then the smaller the attraction of the nucleus on the outer electrons. Even though the valence electrons are being attracted to the nucleus, they are also being repelled by all the inner electrons. The effective nuclear charge causes properties to change that are in the same period (row),. This means that the positive charge of the nucleus is increasing while the number of energy levels is the same. Since the electrons are negatively charged, they are attracted to nuclei, which HN Chemistry Page 3

45 are positively charged. Many of the properties of atoms depend not only on their electron configurations but also on how strongly their outer electrons are attracted to the nucleus. The distance from the nucleus is not changing but the charge of the nucleus is. Atomic radii Atomic radii refers to the size of an atom. It is actually defined as one-half the distance between the nuclei of identical atoms that are bonded together. Atomic radii increases down a group and decreases across a period. It increases toward Fr. The size of the atom naturally increases down a group because the volume of the electron cloud gets bigger as the number of energy levels increase. The size of the atom gets smaller as the atomic number increases because the charge of the positive nucleus is getting larger and attracts the electrons more and more. It is true that the number of electrons is also increasing, but they are all placed in the same energy level, meaning that they are the same distance from the nucleus. The higher the nuclear charge the stronger the pull on those electrons. Task 4b 1. Referring to the periodic table, arrange the following atoms in order of increasing size: P, S, As, Se. 2. Of cesium, Cs, hafnium, Hf, and gold, Au, which element has the smallest atomic radius? Explain your answer in terms of trends in the periodic table. Ionic Radii A positive ion is known as a cation. The formation of a cation by the loss of one or more electrons always leads to a decrease in atomic radius because as electrons are removed the nucleus has a greater attraction for the electrons that remain. The positive ion has lost a whole layer of electrons. A negative ion is known as an anion. The formation of an anion by the addition of one or more electrons always leads to an increase in atomic radius. The extra electrons cause the electron HN Chemistry Page 4

46 cloud to spread out due to greater electron repulsion while the attraction from the protons remains the same. Positive ions are smaller than the atom from which they came, while negative ions are larger than the atom from which they came. Task 4c 1. Distinguish between a cation and an anion. 2. Which of the following cations is least likely to form: Sr 2+, Al 3+, K 2+? 3. Which of the following anions is least likely to form: I -, Cl 2-, O 2-? 4. From each set, determine which atom or ion is the largest. a. Ca, Ca 2+ b. Cl, Cl - c. S, S -, S 2- d. Sr, Sr +, Sr 2+ e. K +, Ca 2+ f. N 3-, O 2-, F - g. Ca 2+, Sr 2+, Ba 2+ h. O 2-, S 2-, Se 2- HN Chemistry Page 5

47 Ionization Energy An electron can be removed from an atom if enough energy is supplied. Any process that results in the formation of an ion is referred to as ionization. Suppose A is any element. A + energy A + + e - A + + energy A 2+ + e - Ionization Energy is the amount of energy required to remove one electron from a neutral atom of an element to make an ion. Specifically, this is the first ionization energy or IE 1. Ionization energy increases toward F. Ionization energy decreases down a group because the farther the valence electron is from the nucleus the lower the attraction between them. Remember that as you move down a group the period number increases, meaning there is another level of electrons being added. Therefore, there will be more shielding electrons between the nucleus and the valence electrons causing repulsion. So, the farther down the element is in the group the lower the ionization energy. Ionization energy increases across a period because the effective nuclear charge (the positive charge) is getting larger while the shielding effect (levels of electrons) remains the same. This means there will be a stronger attraction between the valence electrons and the nucleus, so it will take more energy to remove an electron. HN Chemistry Page 6

48 The second ionization energy is the amount of energy required to remove the second electron from an ion. The third ionization energy is the amount of energy required to remove the third electron from an ion, etc. Successive ionizations require more energy. The more electrons that are removed the harder it is to remove the next. Each successive electron removed from an ion feels an increasingly stronger effective nuclear charge. Notice that there is a big jump in the IE 1 to IE 2 of sodium. Sodium has only one valence electron. It is relatively easy for it to lose that electron when bonding. It will then have a stable configuration with 8 valence electrons. It will be much more difficult to remove the second electron from sodium because this will make the ion become unstable so it takes a lot more energy to do this. Basically, it is easier to remove valence electrons than it is to remove the electron immediately after the last valence electron. Task 4d 1. Referring to the periodic table, arrange the following atoms in order of increasing first ionization energy: Ne, Na, P, Ar, K. 2. In general ionization energy increases toward F. Refer to the graph on ionization energy trends. Considering electron configurations, why do you think B has a lower IE 1 than Be? O has a lower IE 1 than N? 3. Write the equations that show the process for the following. a. The first ionization energy for tin b. The second ionization energy for the tin(i) ion HN Chemistry Page 7

49 4. Explain each of the following. a. Why does Li have a larger first ionization energy than Na? b. The difference between the third and fourth ionization energies of scandium is much larger than the difference between the third and fourth ionization energies of titanium. Why? c. Why does Li have a much larger second ionization energy than Be? 5. Here are the ionization energies for an element in period 2: 900, 1757, 14849, Which element is represented by these energies? Electron Affinity Neutral atoms can also acquire electrons. The energy change that occurs when an electron is acquired by a neutral atom is called the atom s electron affinity. Most atoms release energy when they acquire an electron. Since energy is given off it will have a negative sign (exothermic). A + e - A - + energy Electron affinity increases toward F. It decreases down a group because there is an increase in atomic radius down a group, which decreases electron affinities. Therefore the nucleus is not as likely to gain an electron. Noble gases have a 0 electron affinities. They have a full valence level; therefore they do not gain electrons. Across the periods, the effective nuclear charge is increasing. The shielding effect remains essentially the same. The larger nuclei tend to be more likely to attract electrons other than its own than nuclei that are smaller. Task 4e 1. Order the atoms in each of the following sets from the least electron affinity to the most. a. O, S b. F, Cl, Br, I c. N, O, F d. B, C, N HN Chemistry Page 8

50 Electronegativity Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound. Fluorine is the most electronegative element. It is arbitrarily assigned a value of four. The other values are calculated in relation to this value. Electronegativities increase toward F. Note that the noble gases have a 0 electronegativity. Remember that they are very stable and do not normally form compounds so they do not pull additional electrons toward them. The alkali and alkaline-earth metal are the least electronegative. In compounds, their atoms have a low attraction for electrons. Nitrogen, oxygen, and the halogens are the most electronegative elements. Their atoms attract electrons strongly in compounds. Electronegativities tend to either decrease down a group or stay the same. Task 4f 1. Using the periodic table, place the following element in order of increasing electronegativity: O, Fe, Ge, Sr, S, Zr HN Chemistry Page 9

51 Other Periodic Properties Valence electrons The number of valence electrons can be determined by the group on the periodic table. Notice that the number of valence electrons is related to the group number. # of valence Group electrons (vary) (except He, 2) Oxidation Numbers Oxidation numbers also vary in a regular pattern on the periodic table. Notice that groups 1-13 tend to lose electrons to become positive. Group 14 can either lose or gain 4 electrons. Groups gain electrons to become negative. Group 18 is stable, so it neither loses nor gains electrons. Group # of valence electrons (vary with RN) , HN Chemistry Page 10

52 Topic 5: The Language of Chemistry Chemical Formulas & Chemical Compounds (Chapter 7 in Modern Chemistry) A Chemical Formula Recall that a chemical formula indicates the relative number of atoms of each kind in a chemical compound. For a molecular compound, the chemical formula reveals the number of atoms of each element contained in a single molecule of the compound. C 8 H 18 Unlike a molecular compound, an ionic compound consists of a lattice of positive and negative ions held together by mutual attraction. The chemical formula for an ionic compound represents one formula unit the simplest ratio of the compound s positive ions (cations) and its negative ion (anion). Al 2 (SO 4 ) 3 HN Chemistry Page 1

53 Note how the parentheses are used. They surround the polyatomic anion to identify it as a unit. When there is no subscript written next to an atoms symbol, the value of the subscript is understood to be one. Monatomic Ions By gaining or losing electrons, many main-group elements form ions with noble-gas configurations (recall from Topics 3 & 4). Ions formed from a single atom are known as monatomic ions. Monatomic cations are identified simply by the element s name. Naming monatomic anions is slightly more complicated. First, the ending of the element s name is dropped. Then the ending ide is added to the root name. Examples of Cations Examples of Anions Element Cation Element Cation K K + F F - Potassium Potassium ion Fluorine Fluoride ion Mg Mg 2+ N N 3- Magnesium Magnesium ion Nitrogen Nitride ion The oxidation numbers of main-group monatomic ions can be determined by looking at their group number on the periodic table. The names of many of the ions include Roman numerals. These numerals are part of the Stock system of naming chemical ion and elements. They are used for elements that have more than one oxidation number. Task 5d 1. Label the following as a cation or an anion. a. Te b. W c. Fr 2. Name the following ions. a. Ca 2+ d. N 2+ b. Cu + e. P 3- c. Se 2- f. I - HN Chemistry Page 2

54 Some Common Monatomic Ions Main-group elements lithium Li + beryllium Be 2+ aluminum Al 3+ sodium Na + magnesium Mg 2+ potassium K + calcium Ca 2+ rubidium Rb + strontium Sr 2+ cesium Cs + barium Ba fluoride F - oxide O 2- nitride N 3- chloride Cl - sulfide S 2- phosphide P 3- bromide Br - iodide I - d-block elements and others with multiple ions copper(i) Cu + vanadium(ii) V 2+ vanadium(iii) V 3+ vanadium(iv) V 4+ silver Ag + chromium(ii) Cr 2+ chromium(iii) Cr 3+ tin(iv) Sn 4+ manganese(ii) Mn 2+ iron(iii) Fe 3+ lead(iv) Pb 4+ iron(ii) Fe 2+ cobalt(iii) Co 3+ cobalt(ii) Co 2+ nickel(ii) Ni 2+ copper(ii) Cu 2+ zinc Zn 2+ cadmium Cd 2+ tin(ii) Sn 2+ lead(ii) Pb 2+ HN Chemistry Page 3

55 Binary Ionic Compounds Binary compounds are made of two elements, a metal and a non-metal. In a binary ionic compound, the total numbers of positive charges and negative charges must be equal. In order to write the formula of an ionic compound, you need to write the cation and the anion, then balance the charges so that the sum of the charges equal zero. For example: Mg and Cl: Mg 2+, Cl - There must be 2 Cl s in order for the negative charge to equal and 2 - equal zero. The correct formula for Mg and Cl is MgCl 2. Al and O: Al 3+, O 2- These numbers will not go into each other so I find a number that both of them will go into, 6. I will need 2 Al s to have a charge of 6 + (2 x 3 + = 6 + ). I will need 3 O s to have a charge of 6 - (3 x 2 - = 6 - ). The correct formula for Al and O is Al 2 O 3. Naming Binary Ionic Compounds The nomenclature, or naming system, of binary ionic compounds involves combining the names of the compound s positive and negative ions. To name binary compounds, follow the rules below. 1. Name the cation. (First ion, metal) 2. Name the anion. (Second ion, non-metal) Remember to change the ending of the last ion to ide. For example: KBr is potassium bromide MgCl 2 is magnesium chloride K 2 O is potassium oxide Task 5e 1. Write the formulas for the binary compounds formed between the following elements; a. potassium and iodine d. aluminum and sulfur b. sodium and sulfur e. aluminum and nitrogen c. lithium and phosphorus f. barium and oxygen 2. Name the binary compounds indicated by the following formulas: a. AgCl d. SrF 2 b. ZnO e. CaO c. CaBr 2 f. Ba 2 P 3 HN Chemistry Page 4

56 The Stock System of Nomenclature Some elements, such as iron, form two or more cations with different charges. To distinguish the ions formed by such elements, scientists use the Stock system of nomenclature. This system uses a Roman numeral to indicate an ion s charge. The numeral is enclosed in parentheses and placed immediately after the metal name. Fe 2+ Fe 3+ iron(ii) iron(iii) Names of metals that commonly form only one cation do not include a Roman numeral. Na + Ba 2+ Al 3+ sodium barium aluminum There is no element that commonly forms more than one monatomic anion. Naming a binary ionic compound according to the Stock system is illustrated below: CuCl 2 copper(ii) chloride If you are writing the name you must figure the charge to put with the cation. It is best to start at the back of the compound with the anion, to determine what the oxidation number of the cation will be. Task 5f 1. Write the formula and give the name for the compounds formed between the following ions: a. Cu 2+ and Br - d. Hg 2+ and S 2- b. Fe 2+ and O 2- e. Sn 2+ and F - c. Pb 2+ and Cl - f. Fe 3+ and O 2- HN Chemistry Page 5

57 2. Give the names for the following compounds: a. CuO c. SnI 4 b. CoF 3 d. FeS Compounds Containing Polyatomic Ions First, you must learn your polyatomic ions. Do Not memorize. You will need these all year. Some Common Polyatomic Ions 1 + *ammonium NH * acetate *bromate BrO 3 - *chlorate ClO 3 - * chlorite ClO 2 - CH 3 COO - (C 2 H 3 O - 2 ) *carbonate 2- CO 3 *chromate CrO 4 2- * dichromate Cr 2 O 7 2- *sulfate SO 4 2- *cyanide CN - *sulfite SO 3 2- *hydroxide OH - *iodate IO 3 - *nitrate NO 3 - * nitrite NO 2 - *permanganate MnO 4 - *phosphate PO 4 3- Notice that most of the polyatomic ions are negatively charged and most are oxyanions polyatomic ions that contain oxygen. Some elements can combine with oxygen to form more than one type of oxyanion. Learn all the oxyanions that end in ate and you ll be able to figure the other oxyanions of that element. For example: HN Chemistry Page 6

58 - ClO 4 perchlorate 1 oxygen more than the -ate ion - ClO 3 chlorate the -ate ion - ClO 2 chlorite 1 oxygen less than the -ate ion ClO - hypochlorite 2 oxygens less than the -ate ion Ternary compounds are made up of more an element and a polyatomic ion. To name ternary compounds, you must know your polyatomic ions and follow the rules below. 1. Name the cation. 2. Name the anion. If the anion is an element from the periodic table, change the ending to ide. If the anion is a polyatomic ion, do not change the ending. For example: NH 4 Cl is ammonium chloride CaCO 3 is calcium carbonate K 2 SO 4 is potassium sulfate Writing formulas of ternary compounds Just as in binary ionic compounds, the sum of the positive and negative ions must equal zero. When multiples of a polyatomic ion are present in a compound, the formula for the polyatomic ion is enclosed in parentheses as in aluminum sulfate, Al 2 (SO 4 ) 3. Task 5g 1. Write the formula for the following ternary compounds. a. Lithium nitrate b. Copper(II) sulfate c. Sodium carbonate d. Calcium nitrite e. Potassium perchlorate 2. Give the names for the following compounds. a. Ca(OH) 2 b. KClO 3 HN Chemistry Page 7

59 c. NH 4 OH d. Fe 2 (CrO 4 ) 3 e. KClO Molecular Compounds Molecular binary compounds are made of two non-metals. Unlike ionic compound, molecular compounds are composed of individual covalently bonded units, or molecules. Chemists use two nomenclature systems to name binary molecules. Naming Molecular Compounds with Prefixes The old system of naming molecular compounds is based on the use of prefixes. For example, the molecular compound CCl 4 is named carbon tetrachloride. The prefix tetra- indicates that four chloride atoms are present in a single Numerical Prefixes molecule of the compound. Number The rules for the prefix system of nomenclature of binary molecular compounds are as follows. 1. The element that has the smaller group number is usually given first. If both elements are in the same group, the element whose period number is greater is given first. This element is given a prefix only if it contributes more than one atom to a molecule of the compound. 2. The second element is named by combining (a) a prefix indicating the number of atoms contributed by the element, (b) the root of the name of the element, and (c) the ending -ide. 10 deca- 3. The o or a at the end of a prefix is usually dropped when the word following the prefix begins with another vowel. P 4 O 10 Prefix 1 mono- 2 di- 3 tri- 4 tetra- 5 penta- 6 hexa- 7 hepta- 8 octa- 9 nona- Tetraphosphorus decoxide HN Chemistry Page 8

60 Binary Compounds of Nitrogen and Oxygen Formula N 2 O NO NO 2 N 2 O 3 N 2 O 4 N 2 O 5 Prefix-system name dinitrogen monoxide nitrogen monoxide nitrogen dioxide dinitrogen trioxide dinitrogen tetroxide dinitrogen pentoxide Task 5h 1. Name the following binary molecular compounds using prefixes. a. CI 4 e. As 2 S 3 b. SO 3 f. NCl 3 c. PCl 3 g. SO 2 d. PCl 5 h. ClBr 2. Write the formulas for the following compounds: a. Carbon tetrachloride c. oxygen difluoride b. Dinitrogen trisulfide d. sulfur hexafluoride Naming Molecular Compounds with the Stock System In order to name molecular compounds using the stock system, or for that matter even ionic compounds using the stock system, you have to be able to assign oxidation numbers to the elements in the compound. Namely, you must be able to assign the oxidation number to the first element in the compound. These oxidation numbers, indicate the general distribution of electron among the bonded atoms in a molecular compound or a polyatomic ion. Oxidation numbers are also called oxidation states. A list of rules for assigning oxidation numbers follows. HN Chemistry Page 9

61 Rules for Assigning Oxidation Numbers As a general rule in assigning oxidation numbers, shared electrons are assumed to belong to the more electronegative atom in each bond. 1. The atoms in a pure element have an oxidation number of zero. 2. The more electronegative element in a binary molecular compound is assigned the number equal to the negative charge it would have as an anion. 3. Fluorine has an oxidation number of -1 in all of its compounds because it is the most electronegative element. 4. Oxygen has an oxidation number of -2 in almost all compounds. Exceptions include when it is in peroxides, such as H 2 O 2, in which its oxidation number is -1, and when it is in compounds with fluorine. 5. Hydrogen has an oxidation number of +1 in all compounds containing elements that are more electronegative than it; it has an oxidation number of -1 in compounds with metals. 6. The algebraic sum of the oxidation numbers of all atoms in a neutral compound is equal to zero. 7. The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion. Here are some examples. UF 6 Start with the most electronegative (last), F is -1 (rule 3). Multiply that oxidation number by the number of fluorine atoms. 6 x -1 = -6. Sum must equal zero (rule 6), so U must be +6 H 2 SO 4 Oxygen must have a -2 oxidation number (rule 4). Multiply -2 x 4 = -8 Hydrogen must have a +1 oxidation number (rule 5). +1 x 2 = +2 The sum must equal zero (rule 6), so S must equal +6. ClO 3 - Oxygen must have a -2 oxidation number (rule 4). Multiply -2 x 3 = -6. The sum of a polyatomic ion equals its charge, in this case, -1. That means the oxidation number of chlorine is +5. HN Chemistry Page 10

62 Task 5 i 1. Assign oxidation numbers to each atom in the following compounds or ions. a. HF e. CS 2 b. CI 4 f. H 2 CO 3 - c. H 2 O g. NO 2 2- d. PI 3 h. SO 4 2. Name the following binary molecular compounds using the Stock system. a. CI 4 e. As 2 S 3 b. SO 3 f. NCl 3 c. PCl 3 g. SO 2 d. PCl 5 h. ClBr 3. Write the formulas for the following compounds. a. Sulfur(II) chloride b. Nitrogen(V) oxide c. Carbon(IV) chloride Names & Formulas for Acids An acid is a distinct type of molecular compound. Most acids used in the laboratory can be classified as either binary acids or oxyacids. Binary acids are acids that consist of two elements, hydrogen and a nonmetal. Oxyacids are acids that contain hydrogen, oxygen, and a third element (usually a nonmetal). Oxyacids acids are derived from polyatomic ions. To name a binary acid, follow this guideline. Examples: Hydro STEM ic acid HCl Hydrosulfuric acid hydrochloric acid H 2 S HN Chemistry Page 11

63 To name an oxyacid, follow this guideline. -ate ic, -ite ous -ate ions make ic acids and ite ions make ous acids Examples: H 2 SO 4 H 2 SO 3 HClO 4 sulfuric acid sulfurous acid perchloric acid Nitric acid HNO 3 Nitrous acid HNO 2 Carbonic acid H 2 CO 3 Task 5j 1. Name the following acids. a. H 3 PO 4 c. HI b. HBr d. HC 2 H 3 O 2 2. Write the formulas for the following acids. a. hypochlorous acid c. arsenic acid b. hydroiodic acid d. hydroselenic acid HN Chemistry Page 12

64 Topic 6: Chemical Bonding & Molecular Geometry Chemical Bonding (Chapter 6 in Modern Chemistry) Atoms seldom exist as independent particles in nature. Most substances consist of combinations of atoms that are held together by chemical bonds. A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together. Why do atoms make bonds? It turns out that most atoms are less stable existing by themselves than when they are combined. By bonding with each other, atoms decrease in potential energy, thereby creating more stable arrangements of matter. When atoms bond, their valence electrons are redistributed in way that make the atoms more stable. The way in which the electrons are redistributed determines the type of bonding. In Topic 4 you learned the main-group metals tend to lose electrons to form positive ions, or cations. Nonmetals tend to gain electrons to form negative ions, or anions. Chemical bonding that results from the electrical attraction between cations and anions is called ionic bonding. If a bond is purely ionic, an atom will completely give up electron(s) to another atom. In contrast, atoms joined by covalent bonding share electrons. Covalent bonding results from the sharing of electron pairs between two atoms. In a purely covalent bond, the shared electrons are owned equally by the two bonded atoms. an example of an ionic bond an example of a covalent bond HN Chemistry Page 1

65 Ionic or Covalent? Bonding between atoms is rarely purely ionic or purely covalent. It usually falls somewhere between these two extremes, depending on how strongly the atoms of each element attract electrons. Recall that electronegativity is a measure of an atom s ability to attract electrons. To determine whether a bond is ionic or covalent you have to calculate the difference in electronegativities of the two atoms involved. An electronegativity difference greater than 1.67 is referred to as an ionic bond. Electronegativity differences of 1.67 or less have an ionic bond character of 50% or less. These compounds are typically classified as covalent. Bonding between two atoms of the same element is completely covalent. This is called a nonpolar-covalent bond. This is a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge. Bonds having only 0% to 5% ionic character, or an electronegativity difference equal to or less than 0.3, are considered nonpolar-covalent Ionic Polar -covalent 100% 50% Bonds having an ionic character between 5% and 50%, or with corresponding electronegativity differences of 0.3 to 1.67, are classified as polar-covalent. A polar-covalent bond is a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons. 0.3 Nonpolar-covalent 5% 0 0% Here are some example of determining bond character based on the electronegativities from the periodic table on the next page. Bonding Pair Electronegativity difference Bond type Li and F = 3.00 ionic Cu and S = 0.68 polar-covalent I and Br = 0.30 nonpolar-covalent In summary, subtract the electronegativities, if: Greater than 1.67 ionic Above 0.3 to 1.67 polar-covalent 0.3 or less nonpolar-covalent HN Chemistry Page 2

66 Electronegativities of the elements Atomic radius decreases Ionization energy increases Electronegativity increases Group (vertical) Period (horizontal) 1 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar Lanthanoids Actinoids Periodic table of electronegativity using the Pauling scale Please note, if you do not have the electronegativities, ionic compounds are generally made up of elements that are far apart on the periodic table. For example: K & Cl make an ionic bond. Sr & Br make an ionic bond. S & O make a covalent bond. (Watch out for hydrogen; even though it is in group 1, it has a high electronegativity.) Task 6a K 0.82 Rb 0.82 Ca 1.00 Sr 0.95 Cs Ba Fr 0.7 Ra 0.9 * La 1.1 ** Ac 1.1 Sc 1.36 Y 1.22 Ti 1.54 Zr 1.33 * Hf 1.3 V 1.63 Nb 1.6 Ta 1.5 Cr 1.66 Mo 2.16 W 2.36 Mn 1.55 Tc 1.9 Re 1.9 Fe 1.83 Ru 2.2 Os 2.2 Co 1.88 Rh 2.28 Ir 2.20 Ni 1.91 Pd 2.20 Pt 2.28 Cu 1.90 Ag 1.93 Au 2.54 Zn 1.65 Cd 1.69 Hg 2.00 Ga 1.81 In 1.78 Tl 1.62 Ge 2.01 Sn 1.96 Pb 2.33 As 2.18 Sb 2.05 Bi 2.02 Se 2.55 Br Using the electronegativity chart, determine the type of bond between the atoms in the following compounds. a. AgCl b. K 2 O c. Br 2 d. HCl Te 2.1 Po 2.0 I 2.66 At 2.2 Kr 3.00 Xe 2.60 Rn 2.2 ** Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo Ce 1.12 Th 1.3 Pr 1.13 Pa 1.5 Nd 1.14 U 1.38 Pm 1.13 Np 1.36 Sm 1.17 Eu 1.2 Pu Am Gd 1.2 Cm 1.28 Tb 1.1 Bk 1.3 Dy 1.22 Cf 1.3 Ho 1.23 Es 1.3 Er 1.24 Fm 1.3 Tm 1.25 Md 1.3 Yb 1.1 No 1.3 Lu 1.27 Lr 1.3 HN Chemistry Page 3

67 2. Only using the periodic table, determine if the bonds below are covalent or ionic. Do not use the electronegativity chart. a. XeCl 6 b. CsF c. MgCl 2 d. NO 2 Ionic Bonding and Ionic Compounds An ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. The chemical formula of an ionic compound shows the ratio of ions present in a sample of any size. A formula unit is the simplest collection of atoms from which an ionic compound s formula can be established. For example, one formula unit of sodium chloride, NaCl, is one sodium cation plus one chloride anion. (In the naming of a monatomic anion, the ending of the element s name is replaced with ide.) The ratio of ions in a formula unit depends on the charges of the ions combined. They must achieve electrical neutrality. In other words, the sum of the charges must equal zero. For example: calcium fluoride calcium has a 2 + charge, Ca 2+ fluorine has a 1 - charge, F - In order for their charges to equal zero, there has to be two fluoride ions for every calcium ion. CaF 2 The Formation of Ionic Compounds Consider that a sodium atom and a chlorine atom are approaching each other. The two atoms are neutral and have one and seven valence electrons, respectively. Sodium loses its electron to chlorine forming Na + and Cl -. Sodium atom + Chlorine atom Sodium ion + Chloride ion HN Chemistry Page 4

68 Covalent Bonding Most substances are composed of molecules. A molecule is a neutral group of atoms that are held together by covalent bonds. A molecule may consist of two of more atoms of the same element, as in oxygen, or of two or more different atoms, as in water or sugar. Oxygen molecule, Water molecule, Sucrose molecule, O 2 H 2 O C 12 H 22 O 11 A chemical compound whose simplest units are molecules is called a molecular compound. The composition of a compound is given by its chemical formula. A chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. The chemical formula of a molecular compound is referred to as a molecular formula. A molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound. The molecular formula for water, for example, is H 2 O, which reflects the fact that a single water molecule consists of one oxygen atom joined by separate covalent bonds to two hydrogen atoms. A molecule of oxygen, O 2, is an example of a diatomic molecule. A diatomic molecule is a molecule containing only two atoms. Formation of a Covalent Bond Remember that nature favors chemical bonding because most atoms have lower potential energy when they are bonded to other atoms than they have when they are not bonded. Using the example of two hydrogen atoms, if they are separated by enough distance, they will not influence each other. However, if the two hydrogen atoms approach each other, there will come to a position in which the nucleus of one hydrogen atoms attracts the electron from the other hydrogen atom. The attraction corresponds to a decrease in the total potential energy of the atoms. At the same time the electrons of the two hydrogen atoms are repelling each other. The protons in the nuclei of the two hydrogen atoms are also repelling each other. The repulsion results in an increase in potential energy. The relative strength of attraction and repulsion between the charged particles is dependent on the distance separating the two hydrogen atoms. HN Chemistry Page 5

69 Here the arrows indicate the attractive and repulsive forces between the electrons and nuclei of two hydrogen atoms. Attraction (red) between particles corresponds to a decrease in potential energy of the atoms, while repulsion (blue) corresponds to an increase. The attractive force dominates and continues to pull the two hydrogen atoms closer together until they get to a distance at which the repulsion between like charges equals the attraction between opposite charges. At this position the two hydrogen atoms (now a hydrogen molecule, H 2 ) have their minimum potential energy and are close enough to share electrons. They are now covalently bonded. HN Chemistry Page 6

70 Hydrogen atoms Hydrogen molecule All individual hydrogen atoms contain a single, unpaired electron in a 1s atomic orbital. When two hydrogen atoms form a molecule, they share electrons in a covalent bond. The sharing of the electrons allows each atom to have the stable electron configuration of helium, 1s 2. The tendency is for atoms to achieve a noble-gas configuration by sharing electrons. The Octet Rule Unlike other atoms, the noble-gas atoms exist independently in nature. They possess a minimum of energy existing on their own because of the special stability of their electron configuration. This stability results from the fact that, with the exception of helium and its two electrons in a completely filled outer shell, the noble-gas atoms outer s and p orbitals are completely filled by a total of eight electrons. Other main-group atoms can effectively fill their outermost s and p orbitals with electrons by sharing electrons through covalent bonding. Such bond formations follow the octet rule. Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons init highest occupied energy level. Here is an example of the covalent bonding of hydrogen and fluorine to make hydrofluoric acid, HF. H H F F HN Chemistry Page 7

71 Exceptions to the Octet Rule Most main-group elements tend to form covalent bonds according to the octet rule, However, there are exceptions. Hydrogen forms bonds in which there are only two electrons. Boron, B, and aluminum, Al, has just three valence electrons so it makes bonds surrounded by 6 electrons. Other atoms can be surrounded by more than eight electrons. Examples of elements that make bonds with expanded valences are phosphorus, P, in PF 5 and sulfur, S, in SF 6. Task 6b 1. Use orbital notation to illustrate the bonding in the chlorine molecule, Cl Describe the general location of the electrons in a covalent bond. Metallic Bonding Chemical bonding is different in metals than it is in ionic or molecular compounds. The highest energy levels of most metal atoms are occupied by very few electrons. The properties of metals are due to the highly mobile valence electrons of the atoms that make up a metal. The highest energy levels of most metal atoms are occupied by very few electrons. They have many vacant orbitals. The vacant orbitals in the atoms outer energy levels overlap. This overlapping of orbitals allows the outer electrons of the atoms to roam freely throughout the entire metal. The electrons are delocalized, which means that they do not belong to any one atom but move freely about the metal s network of empty atomic orbitals. These mobile electrons form a sea of electrons around the metal atoms, which are packed together in a crystal lattice. The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons is called metallic bonding. Here are two diagrams to help you understand how the electrons surround the cation in a metallic bond. HN Chemistry Page 8

72 Van der Waals force Van der Waals forces are a group of weak intermolecular forces that vary in strength but are generally weaker than bonds (ionic, covalent, and metallic). We will discuss hydrogen bonding, dipole-dipole interactions, and London dispersion forces (LDF). Hydrogen Bonding Some hydrogen-containing compounds, such as hydrogen fluoride, water, and ammonia, have unusually high boiling points. This is explained by the presence of a particularly strong type of dipole-dipole force. In compounds that contain hydrogen and a very electronegative element (N, O, F) an intermolecular attraction occurs between the molecules. Note that this is not a bond between atoms within the molecule, but an attraction among the molecules. Below, the dotted lines represent the hydrogen bonds between water molecules. This allows water to have relatively high melting and boiling points for hydrogen compounds that bond with group 16 nonmetals. This also explains why ice expands and floats. Dipole-Dipole Attractions A dipole is a molecule or a part of a molecule that contains both positively and negatively charged regions. For example, the molecule, HCl is made with hydrogen and the very electronegative atom chlorine. The electrons in this molecule tend to gather around the chlorine. This makes the hydrogen end more positively charged ( read as partially positive) and the chlorine end more negative ( -, read as partially negative). Each end is a dipole. H Cl + - HN Chemistry Page 9

73 That means that when two hydrogen chloride molecules get close together and they are oriented correctly, the different dipoles will be attracted to each other. London Dispersion Forces There is some type of intermolecular force among all atoms and molecules, even noble gases. These are the weakest of the van der Waals forces and are called the London dispersion forces. London dispersion forces (LDF) are intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles. The more electrons there are the greater the London dispersion forces. So it goes to reason that the bigger the atomic mass or molecular mass the greater the London dispersion forces. That is why boiling points generally increase down a group on the periodic table. Summary of Bonding and their relationship to properties Ionic (Giant lattice) +ve ion and -ve ion formed via transfer of electrons held together in giant lattice with strong electrostatic interaction Metal + Non-Metal ex: KCl, MgF 2, Na 2 SO 4, etc. Strong Bonds high m.p./b.p. poor conductors of electricity when solid (ions not free to move), good when liquid or in solution (dissolved) will dissolve in polar solvents Covalent (Individual molecules) Small groups of atoms covalently bonded together by sharing electrons Non-Metal + Non-Metal ex: CO 2, PCl 5, etc. Strong bonds within molecules, but weak between molecules low m.p/b.p., often liquids or gases at RT; LDF- (induced dipoles) Dipole - perm. dipoles) H-Bonds - (H - N/O/F) poor conductors Metallic (Mixtures of Metals) Close packed array of atoms (ions) with "sea" of free moving electrons Metals ex: Na, Al, Au, Stainless Steel (Fe/C/Cr), Bronze (Cu/Sn), Brass (Cu/Zn) Strong (but flexible) bonds high m.p/b.p. good conductors of electricity ("sea' of electrons) and heat (close packed), malleable (can be shaped), ductile (drawn into thin wires), luster (shiny) HN Chemistry Page 10

74 Task 6c 1. Compound B has lower melting and boiling points than compound A. At the same temperature, compound B vaporizes faster than compound A. If one of these compounds is ionic and the other is molecular, which would you expect to be molecular? Ionic? Explain your reasoning. 2. Analyzing Data. The melting points for the compounds Li 2 S, Rb 2 S, and K 2 S are 900 o C, 530 o C, and 840 o C, respectively. List these three compounds in order of increasing lattice energy. 3. Explain why most metals are malleable and ductile but ionic crystals are not. 4. Explain why metals are good electrical conductors. 5. What is the difference between a formula unit and a molecule. Lewis Structures In Topic 3, you learned about electron dot diagrams for elements. These can be used to represent molecules also. When representing molecules, they are called Lewis structures. Lewis structures are formulas in which atomic symbols represent nuclei and inner-shell electrons, dotpairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons. If only the shared pairs (bonds) are written using dashes, then this will be a structural formula. A structural formula indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule. It is important to learn how to draw the Lewis structures of molecules and to predict the molecular geometry of the molecule. To predict the geometry, you have to consider shared (bonded) electrons and lone (nonbonded) electron pairs surrounding the central atom. To do this, we will use the VSEPR theory. VSEPR stands for valence shell electron pair repulsion. That means that the electron pairs around the atoms (usually the central atom) repel each other to affect the shape of the molecule. VSEPR theory states that the repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible. Unshared electrons repel the most. Remember these molecules are 3-D even though we draw them in 2-D. Drawing Lewis structures HN Chemistry Page 11

75 Here are the steps for writing a Lewis structure for compounds for which you do not know the formula. 1. Calculate the total number of valence electrons taking into account any charges. (add electrons for negative charges and subtract electrons for positive charges) 2. Decide which element is the central atom. Usually this is obvious; if in doubt it will be the least electronegative atom (except for H: it only wants 2 electrons). Put that element in the center and add the other elements around the central atom using lines to represent bonding pairs of electrons. 3. Arrange the remaining electrons to complete the octet of the terminal atoms by placing pairs of dots around the atoms. Place any remaining electrons (dots) on the central atom, if necessary expanding the octet. 4. If the central atom lacks an octet, form multiple bonds (double or triple bonds) by converting non-bonding electrons on terminal atoms into bonding pairs. (Sometimes atoms will remain electron deficient). For example: CCl 4 Valence electrons = 4 + 7(4) = 32 H 2 O Valence electrons = 1(2) + 6 = 8 HN Chemistry Page 12

76 Now practice with your teacher the following: NH 3 HF PF 5 NH 4 + PCl 6 - CO 3 2- HN Chemistry Page 13

77 Task 6d 1. Draw a Lewis structure for each of the following molecules. You will come back to this section for later tasks, so be sure to be neat and orderly. a. SCl 2 b. PI 3 c. Cl 2 O d. NH 2 Cl e. SiCl 3 Br f. ONCl g. SO 4 2- HN Chemistry Page 14

78 h. ClO 2 - i. BeCl 2 You can also use VSEPR theory to predict the shapes of a molecule. Here is a link that might help you understand the shapes. You will need to learn the chart on the next page. HN Chemistry Page 15

79 HN Chemistry Page 16

80 Task 6e 1. Using the Lewis structures that you drew in 6d, determine the molecular geometry for each. Hybridization VSEPR theory helps determine the shapes of a molecule but it does not reveal the relationship between a molecule s geometry and the orbitals occupied by its bonding electrons. To explain how the orbitals of an atom become rearranged when the atom forms covalent bonds, a different model is used. This model is called hybridization, which is the mixing of two or more atomic orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal energies. Methane, CH 4, provides a good example of how hybridization is used to explain the geometry of molecular orbitals. The orbital notation for carbon s valence electrons is 2s 2 2p 2. We know form experiments that a methane molecule has tetrahedral geometry. How does carbon form four equivalent, tetrahedrally arranged covalent bonds? The one s orbital and the three p orbital hybridize to form four new, identical orbitals called sp 3 orbitals. 2p sp 3 2s Carbon s orbitals Before hybridization Carbon s orbitals after sp 3 hybridization Hybrid orbitals are orbitals of equal energy produced by the combination of two or more orbitals on the same atom. The number of hybrid orbitals produced equals the number of orbitals that have combined. A molecule with 3 bonding areas will be sp 2, while a molecule with 4 bonding areas will be sp 3. Trigonal bipyramidal molecules will be dsp 3 and octahedral molecules will be d 2 sp 3. HN Chemistry Page 17

81 Polarity & Dipole forces The strongest intermolecular forces exist between polar molecules. Polar molecules act as tiny dipoles because of their uneven charge distribution. A dipole is created by equal but opposite charges that are separated by a short distance. The direction of the dipole is from the dipole s positive pole to its negative pole. A dipole is represented by an arrow with a head pointing toward the negative pole and a crossed tail situated at the positive pole. The negative pole will be the atom that is the most electronegative. For example, hydrochloric acid, HCl: H Cl This is a polar bond and a polar molecule because the charges are unevenly distributed. If the charge distribution is evenly distributed, the molecule will be nonpolar. For example CH 4 : H H C H Notice that all the dipole point toward the C. Here the charge is evenly distributed. Therefore the molecule is nonpolar even though the individual bonds are polar. H Task 6f 1. Go back to Task 6d and label the hybridization and the polarity of each molecule. HN Chemistry Page 18

82 Topic 7: The Mole Concept Relating Mass to Numbers of Atoms (Chapter 3 in Modern Chemistry beginning on p.82) In order to understand the quantitative parts of chemistry, there are three very important concepts the mole, Avogadro s number, and molar mass. These provide the basis for relating masses in grams to number of atoms. The Mole The mole is the SI unit for amount of substance. A mole (abbreviated mol) is the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12. The mole is a counting unit, just like a dozen is. Avogadro s number The number of particles in a mole has been experimentally determined in a number of ways. The best modern value is x This is the number of particles (atoms, ions, molecules, formula units) in a mole. It is called Avogadro s number. Avogadro s number x is the number of particles in exactly one mole of a pure substance. For most purposes, Avogadro s number is rounded to x Watch this video for an idea of how many particles Avogadro s number really is. Molar Mass An alternative definition of mole is the amount of a substance that contains Avogadro s number of particles. The mass of one mole of a pure substance is called the molar mass of that substance. Molar mass is usually written in units of g/mol. The molar mass of an element is numerically equal to the atomic mass of the element in atomic mass units (which can be found in the periodic table). Using Chemical Formulas (Chapter 7 in Modern Chemistry beginning on p.237) As you have seen, a chemical formula indicates the elements as well as the relative number of atoms or ions of each element present in a compound. Chemical formulas also allow chemists to calculate a number of characteristic values for a given compound. In this section, you will learn how to use chemical formulas to calculate the molar mass, formula mass, and the percentage composition by mass of a compound. HN Chemistry Page 1