Matter What is Chemistry? Chemistry is the study of matter and the changes it undergoes. What is matter? Matter is anything that has mass and occupies space. Chemists use a scientific method to study matter. Hypothesis (tentative explanation or prediction) Experiment (controlled procedure) Data (quantitative and qualitative) Theory (provides explanations and predictions) Scientific Laws (mathematical equations/models) Classification of Matter
Phases of Matter Properties of Matter Chemical Property Any characteristic that can be determined only by changing a substance's molecular structure. Physical Property Any characteristic that can be determined without changing the substance's chemical identity. Extensive Property Any characteristic of matter that depends on the amount of matter being measured. Intensive Property Any characteristic of matter that does not depend on the amount of the substance present Units of Measurements Many properties of matter are quantitative ; that is, they are associated with number. When a number represents a measured quantity, the units of that quantity must always be specified.
SI Units Physical Quantity Name of Unit Abbreviation Mass Kilogram Kg Length Meter m Time Second s Temperature Kelvin K Amount of Substance Mole mol Electric current Ampere A In the metric system, prefixes are used to indicate the decimal fractions or multiples of various units. Prefixes used in Metric System Prefix Abbreviation Meaning Giga G 10 9 Mega M 10 6 Kilo k 10 3 Deci d 10 1 Centi c 10 2 Milli m 10 3 Micro μ 10 6 Nano n 10 9 Pico p 10 12 Temperature: Temperature is a measure of hotness or coldness of an object (kinetic energy) The temperature scales commonly used in scientific studies are Celsius and Kelvin. K = 273 + o C
Derived SI Units Volume: cm 3 = cc= ml Density: D= m/v Uncertainty in Measurements Uncertainties always exist in measured quantities. Measured quantities are generally reported in such a way that only the last digit is uncertain. Precision and accuracy are two closely related terms with uncertainties of measurements. Accuracy refers to how close to the true value a given measurement is. Precision refers to how well a number of independent measurements agree with one another. To indicate the uncertainty in a measurement, the value you record should use all the digits you are sure of, plus one additional digit that you estimate. Significant Figures The total number of digits in a measurement is called the significant figures. More number of significant figures in a measurement less uncertain the measurement is. RULES: Non zero integers always count as significant figures. Ex. 234 or 1768 Zeros present in a measured quantity present problems (to find the number of significant figures) because they can be used in two ways: as part of the measured value or to position a decimal point. Zeros between two other significant digits (enclosed zeros or captive zeros) are always significant. Ex. 1107 or 50.001 Zeros at the beginning (leading zeros) of a number are never significant. Ex. 0.0025 or 0.02 Zeros at the end (trailing zeros) of a number are significant if they involve a decimal point. They are the zeros at the right end of the number. Ex. 0.0500 or 3.00 or 100. A problem arises when a number ends with zeros but contains no decimal point. for example, the number of significant figures in 10300 g can be obtained without confusion if it is written in scientific notation.
Significant Figures in Calculations In Multiplication and Division the result must be reported with the same number of significant figures as the measurement with the fewest significant figures. When the result contains more than the correct number of significant figures, it must be rounded. Ex. 0.725 cm x 0.51 In Addition and Subtraction, the limiting term is the one with the smallest number of decimal places. Ex. 49.146 + 72.13 The Atom Dalton s Atomic Theory Dalton s Atomic Theory of matter involved the following assumptions: 1. Each element is composed of extremely small particles called atoms. 2. All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all the other elements. 3. Atoms are neither created nor destroyed in chemical reactions. 4. Compounds are formed when atoms of more than one element combine in a simple numerical ratio. Law of Conservation of Mass Matter is not created or destroyed only changed (Lavoisier) Using balances, he found that the total mass does not change after chemical reactions. Law of Definite Proportions In a given chemical compound, the elements are always combined in the same proportions by mass. (Proust) Each compound has a definite percent of each element in that compound. The Structure of the Atom Atom Word is from Greek, meaning unable to be cut or indivisible. Proposed by Greek philosophers: Democritus, Empedocles (about 450 BC) A philosophical idea, rather than scientific theory. Basically speculative, not based on observation or measurement. Some of the Greek philosophers ascribed shapes to the atoms. Atoms of water: spherical (smooth, flow) Atoms of earth: cubical (rigid, solid) Atoms of fire: jagged (hurts to touch)
Modern chemistry and modern atomic theory, relating to scientific observation and measurement, dates back approximately 200 years, to the late 18th and early 19th centuries. Subatomic Particles what atoms are made up of Particle Symbol Mass (g) Charge (Coulomb) Relative Charge Location in Atom Proton p + 1.67262x10 24 +1.6022x10 19 +1 Inside Nucleus Neutron n 1.67262x10 24 0 0 Inside Nucleus Electron e 9.10939x10 28 1.6022x10 19 1 Outside Nucleus Atomic Number, Mass Number, and Nucleus atomic number (Z) = t he number of protons in the nucleus of an atom mass number (A) = the sum of the number of protons and neutrons in an atom element symbol (X) = the symbol of the element Note: mass number= number of protons + number of neutrons Therefore. mass number = atomic number + number of neutrons. A= Z + number of neutrons.. Number of neutrons = A Z Note: For any given element on the periodic table: Number of protons = Number of electrons In order to symbolically represent elements and isotopes chemists use the following notation: Mass Number X Atomic number The nucleus, containing protons and neutrons, accounts for almost all the mass of the atom.
The electron cloud accounts for almost all the space (volume) of the atom, but almost none of the mass. Isotopes are atoms with identical atomic numbers but different mass numbers (that is, same number of protons but different number of neutrons) A given element, with a fixed atomic number, will have several isotopes with different mass number. Such isotopes differ only in the number of neutrons in the nucleus. Thus, unlike Dalton s original assumption, the same element can have atoms with different masses. Writing Ionic Compounds How to ions combine to form neutral compounds (Ionic Compounds)? * A chemical compound must have a net charge of zero. * There must be both positive and negative ions present * the number of cations and anions must be such that the net charge is zero. *use criss cross method Naming Compounds (Chemical Nomenclature) Naming is based on the division of chemical compounds into categories. The major division is between organic and inorganic compounds. Organic Compounds contain carbon, usually in combination with hydrogen, oxygen, nitrogen, or sulfur. Inorganic Compounds all others Among inorganic compounds we will consider three categories: ionic compounds, molecular compounds, and acids. Ionic Compounds Rules: 1. Positive Ions (cations) Always Named First a. (Type 1) Monatomic cations formed from metal have the same name as the metal. b. (Type 2) If a metal can form different cations, ex. transition metal ions and some group 14, the positive charge is indicated by a roman numeral in parenthesis following the name of metal 2. Negative Ions (anions) Always Named Second
a. The names of monatomic anions are formed by replacing the ending of the name of the element with ide NOTE: The number of positive charges on each of the cations in the same as the group number (1A, 2A, 3A). The number of the negative charge on the anion is equal to the number of spaces to the right that we have to move in the periodic table to get to a noble gas (4A, 5A, 6A, 7A). Ionic compounds are electrically neutral. In the chemical compound, the subscripts in the formula must produce an electrically neutral formula unit as shown below. Polyatomic Ions You should know the number of atoms, the charge, and the name for the polyatomic ions. Use this Quizlet Names and Formulas of Binary Molecular Compounds(Covalent)
Rules: 1. The first element in the formula is named first! 2. The second element is named as though it were an anion with ide ending 3. Greek Prefixes are used to indicate the number of atoms of each element. The prefix mono is never used in the first element. When the prefix end in a or o and the name of the second element begins with a vowel (such as oxide), the a and o of the prefix is often dropped. NOTE: The name of the element father to the left in the periodic table is usually written first. An exception to the rule occurs in the case of the compounds that contain oxygen. Greek Pre fixes Number Indicated Prefix 1 Mono 2 Di 3 Tri 4 Tetra 5 Penta 6 Hexa 7 Hepta 8 Octa 9 Nona 10 Deca Naming Acids Acids are substances which give off H + ions. We name the acids based on the negative ion which is left over after the H + ion is given off. Negative ion suffix ate ite ide Corresponding Acid ic acid ous acid hydro ic acid Examples
H 2 SO 4 (sulfate ion) Sulfuric acid H 2 SO 3 (sulfite ion) Sulfurous acid HNO 3 (nitrate ion) Nitric acid HNO 2 (nitrite ion) Nitrous acid H 3 PO 4 (phosphate ion) Phosphoric acid HClO 4 (perchlorate ion) Perchloric acid HClO 3 (chlorate ion) Chloric acid HClO 2 (chlorite ion) Chlorous acid HClO (hypochlorite ion) Hypochlorous acid HCl (chloride ion) Hydrochloric acid HF (fluoride ion) Hydrofluoric acid Naming Bases On the simplest level, bases are substances which give off hydroxide ion, OH. bases are hydroxides of group I and group II. The most common Sodium hydroxide Potassium hydroxide Calcium hydroxide NaOH KOH Ca(OH)2 Acids and bases react with each by a process called neutralization. The H + ion from the acid reacts with the OH ion from the base forming water. Example: Reaction of hydrochloric acid and sodium hydroxide.. HCl + NaOH NaCl + HOH In a future lecture,we will study these reactions in more detail. Naming Hydrates In a hydrate, each formula unit of the compound has associated with it a certain number of water molecules. Mg 2 Cl 2 6 H 2 O LiC 2 H 3 O 2 2 H 2 O Atomic Mass The atomic mass (atomic weight) of an element is the average of the isotopic masses, weighted according to the naturally occurring abundances of the isotopes of the element.
The Mole Concept Consider the following approximate atomic masses of some elements. These numbers represent the relative mass of each element. The units are atomic mass units (abbreviated amu, and given the symbol u), a very small unit of mass. H He C O Cu Pb U 1.0 u 4.0 u 12.0 u 16.0 u 63.5 u 207 u 238 u This scale tells us relative masses, with carbon 12 being defined as exactly 12 u. A helium atom has a mass 4 times that of hydrogen. A carbon atom has a mass 3 times that of helium (ratio of 12/4). A uranium atom is almost 20 times heavier than a carbon atom (ratio of 238/12). Now suppose we have a number of atoms (called an amount ) that has a mass equal to the atomic masses in grams. Let s call this amount N atoms. N atoms of hydrogen weigh 1.0 grams This reasoning should be clear: The same N atoms (the same amount) of each element would have a mass equal to its atomic mass in grams. Mass N atoms of hydrogen 1.0 g N atoms of helium 4.0 g N atoms of carbon 12.0 g N atoms of oxygen 16.0 g N atoms of copper 63.5 g N atoms of lead 207 g N atoms of uranium 238 g This amount, the number of atoms that has a mass equal to the atomic mass in grams, is called a mole. A mole of an element always has the same amount of atoms, so far designated by N. But a mole does not have to refer only to atoms. We can have compounds in which a molecule (a group of atoms) is the smallest unit.
On the atomic mass scale: Mass 1 molecule H 2 O 18.0 u 1 molecule CO 2 44.0 u 1 molecule C 6 H 12 O 6 180 u Here we find the mass of each molecule ( molar mass ) by just adding up the atomic masses in the formula. By the same reasoning illustrated above: Mass N atoms of carbon 12.0 g N molecules of H 2 O 18.0 g N molecules of CO 2 44.0 g N molecules of C 6 H 12 O 6 180 g N represents the number of elementary units in a mole. Definition of A Mole The mole is the SI unit of amount. An atom of carbon 12 has a mass of exactly 12 u. (this defines the amu scale) A mole of carbon 12 is the amount of carbon 12 having a mass of exactly 12 grams. A mole of any element contains the same number of atoms as is contained in exactly 12 grams of carbon 12. A mole of any substance contains the same number of elementary entities as is contained in exactly 12 grams of carbon 12. Elementary entities may be atoms, molecules, ions, electrons, or any other appropriate unit. Working definition: A mole of any substance contains the number of elementary units (atoms or molecules) which has a mass equal to its atomic or molecular mass in grams.
So far, we have designated that amount by the symbol N. The value of N has been determined experimentally. It is known as Avogadro s number. N = 6.022 x 10 23 particles (four significant figures) This is a very large number. To get some idea of its magnitude, consider this example. Chemical Equations Chemical reactions are symbolically represented by chemical reactions. Should Include: compounds with correct formulas separate the reactants and products with arrow indicate phases of matter balance the equation Example: hydrogen gas reacts with oxygen gas in the air to form water. Basic Types of Reactions: Combustion reaction: Compound + O 2 Product Synthesis (or Combination) reaction: A + B AB Decomposition reaction: AB A + B Single Displacement: AB + C AC + B Double Displacement: AB + CD AD + CB