In 1808, John Dalton introduced the idea of atoms and supported it with experiments.

Transcription

1 Chemistry 11 Unit VIII Atomic Theory and the Periodic Table Notes Development of the Atom Dalton s Atomic Theory In 1808, John Dalton introduced the idea of atoms and supported it with experiments. Dalton s Atomic Theory is as follows 1. Elements are made up of extremely small particles called atoms. 2. The atoms making up a particular element are identical and different types of atoms have different properties. 3. Each chemical compound is unique and consists of a particular combination of specific types of atoms put together in a distinctive way. 4. Chemical reactions involve the reshuffling of the atoms. Products are made form the same atoms as the reactants. Thomson Model of the Atom 1897 J.J Thompson discovered that atoms contain negatively charged particles (later named electrons). He showed that positively charged particles could also be obtained from an atom. Because atoms are neutral, Thompson proposed that atoms must possess both positive and negative charges. Thompson did not how they were arranged, but proposed that an atom consisted of a ball of positive particles with the negative particles distributed throughout. Think chocolate chips, or Plum Pudding The Rutherford Model of the Atom (Gold Foil Experiment) Rutherford s Gold Foil experiment was ground breaking. Rutherford was a student of J.J Thompson and firmly believed in the plum pudding model. Rutherford fired alpha particles (radioactive) at gold foil. Some of the particles were deflected and others bounced straight back, This was amazing, as it showed that there had to be something in the atom that could stop an alpha particle. If the plum pudding model was right, this observation would not have been seen. Simmer Down Inc. (Updated Dec 2012) 1

2 Rutherford proposed that the atom consisted of a tiny, positively charged nucleus surrounded by a cloud of negatively charged electrons. The nucleus contains almost all the mass and consists of protons and neutrons. The number of electrons surrounding the nucleus equals the number of protons, thus making it neutral. This model explained the charged nature of an atom, but did not account for all its mass. Rutherford believed a neutral particle also existed that added mass, but no charge, Atom consists of a tiny, positively-charged nucleus surrounded by a cloud of negativelycharged electrons. The nucleus contains almost all of the mass of the atom and consists of protons and neutrons. The number of electrons surrounding the nucleus equals the number of protons in the nucleus, so as to make the atom electrically neutral J. Chadwick discovered the neutron as predicted by Rutherford The following table summarizes the properties of the 3 major atomic particles. The Properties of Atomic Particles Particle Symbol Charge Molar Mass (g) Found in Electron e x Outside nucleus Proton P x Inside nucleus Neutron N x Inside nucleus Rutherford s work was more concerned with the atomic nucleus and not with what the electrons were doing. This planetary model of electrons suggested that the electrons whirl around the nucleus. (Same way planets resolve around the sun). Niels Bohr 1913 Bohr came up with an equation to explain the way electrons behave. He proposed that electrons only exist in specific energy states. These energy states are associated with specific circular orbits, which the electron can occupy around the atom. Simmer Down Inc. (Updated Dec 2012) 2

3 When the electron absorbs or emits a specific amount of energy it instantaneously moves from one orbit to another. The greater the energy, the further the orbit from the atomic nucleus. Bohr proposed that the electrons in an atom are restricted to having certain specific energies and are restricted to following specific paths called orbits at a fixed distance from the nucleus. Electrons were only allowed to emit or absorb energy when they moved from one orbit to another. Bohr model of the Atom Bohr proposed that the electrons in an atom are restricted to having certain specific energies and are restricted to following specific paths called orbits at a fixed distance from the nucleus. Electrons were only allowed to emit or absorb energy when they moved from one orbit to another. Bohr model ran into some problems as well. The notion of electrons orbiting a nucleus along a specific path had to be abandoned and replaced by the idea that different electrons, depending on their energies, simply occupy particle regions of space called orbitals. Overall picture requires the use of both models and thus is referred to The Rutherford-Bohr model. Atomic Number and Atomic Mass Elements are differentiated from one another by the number of protons in the nucleus. Example: H has 1 proton in its nucleus He has 2 protons in its nucleus Na has 11 protons in its nucleus The Atomic Number of an atom is the number of protons in the nucleus. A neutral atom has no overall net charge, thus the number of protons must equal the number of electrons. In a Neutral atom Here are 3 things that you need to remember. They are usually on any standard periodic table. Number of Protons = Number of Electron Simmer Down Inc. (Updated Dec 2012) 3

4 If electrons are added or subtracted from a neutral atom, the resulting particle is called an ion. Adding a Negative charge (electron) produces a Negative ion (Anion). Taking away an Electron produces a Positive ion (Cation). Mass Number of an atom is the total number of protons and neutrons. Since Atomic Number = The # of protons Atomic mass (Mass number) = # of protons + # of neutrons Number of neutrons = Mass Number Atomic Number The mass number and atomic mass (weight) is NOT the same thing. The atomic mass is an average of the naturally occurring isotopes, hence the decimal values. The mass number is usually calculated by rounding the atomic mass to the nearest whole number ***For the sake of consistency, use the periodic table that has the mass numbers shown*** Examples: 1. How many protons, electrons and neutrons does Ca contain? 2. How many protons, electrons and neutrons does Ga +3 contain? 3. How many protons, electrons and neutrons does P -3 contain? 4. How many protons, electrons and neutrons does 224 Ra +2 contain? Isotopes Isotopes are atomic species having the same atomic number (protons) but different atomic masses (neutrons). There are 3 different types of Hydrogen, each of which has a special name. 1 1 H = H = Ordinary hydrogen 2 1 H = D = Deuterium (sometimes called heavy hydrogen) 3 1 H = T = Tritium (sometimes called radioactive Hydrogen) ***No other isotopes have a special name*** Simmer Down Inc. (Updated Dec 2012) 4

5 Natural Mixtures of Isotopes The molar mass of Chlorine is 35.5 g. Since there can not be 0.5 of a proton or neutron, then 35.5 g must represent an average value of a mixture of isotopes. Similar to saying that the typical Canadian family has 2.5 children this refers to an average. How else could you get half a child You need to be able to calculate the average atomic mass for isotopic data. Examples: 1. Chlorine exists as a mixture of 75.77% Cl-35 and 24.23% Cl-37. If the precise molar mass of Cl-35 is g / mol and Cl-37 is g / mol, what is the average molar mass of the chlorine atoms? Solution: mass of Cl-35 = (0.7577) x ( g / mol ) = g / mol mass of Cl-37 = (0.2423) x ( g / mol ) = g / mol total mass = g g = g / mol *If the exact masses of the isotopes are not given in the question, the atomic masses can be used instead.* 35 Cl = 75.77% and 37 Cl = 24.23% Average mass = ( x 35) + ( x 37) = g / mol 2. Calculate the average atomic mass of the following: 95% 14 N, 3% 15 N, 2% 16 N Electronic Structure of the Atom (This is where is gets fun ) Remember Niels Bohr proposed that electrons only exist in specific energy states. When an electron absorbs or emits a specific amount of energy it instantaneously moves from one orbital to another. An energy level is a specific amount of energy, which an electron in an atom can possess. ( n is the number of the energy level. AKA = principal quantum number) The energy difference between two particular energy levels is called the quantum of energy associated with the transition between the two levels. An orbital is the actual region of space occupied by an electron in a particular energy level. Simmer Down Inc. (Updated Dec 2012) 5

6 Orbitals come in different shapes and sizes. We use letters to distinguish between them. They are: s, p, d and f. The use of an energy level (n) and an orbital make up the first part of the electron configuration. Example:. 1s, 2p, 4d etc. ***Note There is a specific order in how the orbitals fill, it is NOT random*** A Shell is the set of all orbitals having the same n-value (Ex. 3s, 3p. 3d) A Subshell is a set of orbitals of the same type. For example, the set of five 3dorbitals in the 3 rd shell is a subshell. The Rules that governing which types of orbitals can occur for a given energy level, and how many orbitals of a given type can exist, are given below 1. For a given value of n, n different types of orbitals are possible i. For n = 1: only the s-type is possible ii. For n = 2: the s- and p- types are possible iii. For n =3: the s-, p- and d-types are possible iv. For n = 4: the s-, p-, d- and f- types are possible 2. An s-type subshell consists of ONE s-orbital A p-type subshell consists of THREE p-orbitals A d-type subshell consists of FIVE d-orbitals An f-type subshell consists of SEVEN f-orbitals Simmer Down Inc. (Updated Dec 2012) 6

7 Please look at your energy level diagram handout. The addition of electrons to the orbitals of an atom follow two rules 1. As the atomic number increases, electrons are added to the orbitals. To ensure the lowest possible energy for the atom, electrons are added to the orbitals having the lowest energy first. (n = 1 1 st, then n = 2 2 nd etc) 2. A maximum of 2 electrons can be placed in each orbital. This means that there can be a maximum of: 2 electrons in a s-type subshell 6 electrons in a p-type subshell 10 electrons in a d-type subshell 14 electrons in an f-type subshell Just remember, each dash line on that handout means 2 ELECTRONS Writing Electronic Configurations for Neutral Atoms An electron configuration is a description of which orbitals in an atom contain electrons and how many electrons are in each orbitals. Remember to always start with the lowest energy level first. The number of electrons in a neutral atom is the same as the atomic number. First step is to figure out how many electrons you have (neutral atom = atomic number), then start at the lowest energy level (1s) and just keep adding till you have none left. The addition of electrons to the orbitals of an atom follows 2 simple rules: i) As the atomic number increases, electrons are added to the available orbitals. Electrons are added to the orbitals having the LOWEST energy first. The order in which orbitals are filled is 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p ii) A maximum of 2 electrons can be placed in each orbital. This means there can be a MAXIMUM of: 2 electrons in a s-type subshell 6 electrons in a p-type subshell 10 electrons in a d-type subshell 14 electrons in an f-type subshell This space for rent Simmer Down Inc. (Updated Dec 2012) 7

8 The electron configuration of most elements can be easily determined by using the orbital version of the periodic table shown below. Examples: 1. Write the electron configuration of Na Determine how many electrons. Na is neutral, thus protons = electrons Atomic number = 11 (# of protons), thus we have 11 electrons. Start at 1s and just keep going 1s 2 2s 2 2p 6 3s 1 2. Write the electron configuration of Zinc 3. Write the electron configuration of Strontium Core Notation The set of electrons belonging to a given atom can be divided into two subsets: the core electrons and the outer electrons. The CORE of an atom is the set of electrons with the configuration of the previous noble gas (He, Ne, Ar, Kr, etc) having an atomic number less than that of the atom being considered. Simmer Down Inc. (Updated Dec 2012) 8

9 The OUTER electrons consist of all electrons outside the core. Since core electrons normally take part in chemical reactions, they are not always explicitly included when writing the electron configuration. Core notation is a way of showing the electron configuration in terms of the core and the outer electrons. Examples: The rules are core notation is as follows 1. Locate the atom and note the noble gas at the end of the row ABOVE the element. 2. Start to write the electron configuration as usual, but replace the part of the electron configuration corresponding to the configuration of the noble gas with the symbol for the noble gas in square brackets. Follow the core symbol with the electron configuration of the remaining outer electrons. 1. Write the electron configuration, using core notation, for Sulphur (S) and Rubidium (Rb) 2. Write the electron configuration, using core notation, for Nickel (Ni). Electron Configurations Exceptions There are two notable exceptions to the electron configurations. These are Chromium (Cr) and Copper (Cu). You will need to MEMORIZE them. Cr = 24 e - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Cu = 29 e - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 Simmer Down Inc. (Updated Dec 2012) 9

10 Writing Electron Configurations for Ions To write the electron configuration of a negative ion, add electrons (equal to the charge) to the last unfilled subshell, starting where the neutral atom left off. Examples: Write the electron configuration for S -2 Sulphur (S) has 16 electrons, so S -2 should have 18 electrons. Thus, just start from the bottom (1s) and add till you run out. 1s 2 2s 2 2p 6 3s 2 3p 6 Write the electron configuration for N -3 The electron configuration of positive ions involves two rules. 1. Starting with the neutral configuration in core notation, remove electrons from the outermost shell (largest n-value) first 2. If there are electrons in both the s- and p- orbitals of the outermost shell, the electrons in the p-orbitals are removed first. Summarized as follows Write the core notation for the atom; remove electrons in the following order p-electrons before s-electrons before d-electrons Examples: 1. Write the electron configuration for Ca Write the electron configuration for Sn Write the electron configuration for Ge +4 Simmer Down Inc. (Updated Dec 2012) 10

11 Predicting the number of Valence Electrons Valence electrons are the electrons, which can take part chemical reactions. In other words Valence electrons are all the electrons in an atom EXCEPT those in the Core, or In the filled d- or f- subshells Examples: Al = [Ne] 3s 2 3p 1 Pb = [Xe] 6s 2 4f 14 5d 10 6p 2 Xe = [Kr] 5s 2 4d 10 5p 6 OR [Xe] has 3 valence electrons has 4 valence electrons has 0 valence electrons (noble gas configuration) How many valence electrons in Calcium? The Periodic Table Chemists were starting to discover a lot of elements. 52 by 1817 and that number rose to 62 by Some kind of organization was needed, but 1817 was too early. Chemists needed to refine their skills. First attempt was done in the 1820 s, but the real fun started to happen in 1850 s and 1860 s William Odling separated the known elements into 13 groups based on their physical and chemical properties. (It was a start) Between 1863 and 1866 John Newlands showed that by assigning Hydrogen an arbitrary mass of 1 and ordering the known elements by their masses, every eighth element shared a common set of properties. He called this the law of octaves Newland s work was good, but failed to predict elements. Plus, the guy could not make up his mind. He kept changing the way he ordered elements would be the year of a breakthrough. Russian Chemist, Dimitri Mendeleev published a method of organizing the elements according to both their masses and their properties. Mendeleev showed that when the elements are listed according to masses, certain properties recur PERIODICALLY. He broke the list into a series of rows and columns. He called each horizontal row a PERIOD and a vertical column a GROUP. Simmer Down Inc. (Updated Dec 2012) 11

12 This is what the early version looked like He placed elements in certain groups based on their properties in spite of contrary indications by its mass (brilliant ). Mendeleev left gaps in his table for elements, which he proposed had yet to be discovered. He was able to predict the properties and characteristics of the undiscovered elements so accurately that when they were discovered the predicted value was almost bang on. The periodic table allowed chemists to organize and understand their data and predict new properties. Modern Periodic Table The modern periodic table is organized according to atomic number rather than atomic mass. This solved the problems where different isotopic abundances caused the masses to be out of order. (Example Ar and K, Co and Ni) The periodic law summarizes the periodic table. The Periodic Law: The properties of the chemical elements recur periodically when the elements are arranged from lowest to highest atomic numbers Major Divisions in the Periodic Table Period: The set of all elements in a given row going across the table. Group or Family: The set of all elements in a given column going down the table. There are several special groups, rows or blocks of elements. They are: Alkali Metal = elements in the first column (Except Hydrogen) Alkaline Earth Metals = The elements in the second column Simmer Down Inc. (Updated Dec 2012) 12

13 Halogens = second column from the end on the right hand side. Headed by Fluorine. Noble Gases = Far right side of the table. Headed by helium. Lanthanides = elements in the first row shown underneath the table. Starting with lanthanum. Actinides = underneath the Lanthanides. Starting with actinium. Metals, Non-Metals and Semiconductors The elements can also be classified based on their metallic characteristics. Properties of Metals Reflect light when polished. Shiny and have a metallic lustre. Are opaque. Good conductors of electricity and heat. Generally, but not always, flexible when in sheets Malleable (Hammered or rolled into thin sheets) Ductile (Drawn into wires) Usually solid at room temperature (Mercury is an exception) Properties of Non-Metals Gases, liquids or brittle solids at room temperature. Poor conductors of heat and electricity Solids are dull to lustrous in appearance and opaque to translucent. Non Metals can be divided into two subgroups: Non-metals with very low electrical conductivities Non-metals with fair to moderate conductivity. Semiconductor: A non-metal having an electrical conductivity, which increases with temperature. Simmer Down Inc. (Updated Dec 2012) 13

14 Semiconductors (Metalloids or Semimetals) have properties which resemble metals more than non-metals Important difference is that metal conductivity decreases with increasing temperature whereas the electrical conductivity of semiconductors increases with increasing temperature. See the figure below for the divisions in the periodic table Periodic Trends There are several trends that you must be able to describe and draw. The trends you will need to know are as follows: Metallic Properties 1. Metallic properties 2. Atomic Radius 3. Ionization energy 4. Electronegativity The properties of the elements change from metallic to non-metallic going from left to right across the table. Elements become more metallic (or better metals) going down a family in the periodic table. Atomic Radius Atomic radii of an atom decreases going across a row left to right. Atomic radii of an atom increases going down a group. Why? Going from left to right across a given period, the atomic number (and the number of protons) increases, and the positive charge on the nucleus increases. Simmer Down Inc. (Updated Dec 2012) 14

15 This increase in atomic number also brings an increase in the number of electrons surrounding the nucleus. All the electrons in a given shell can be assumed to have the same average distance from the nucleus. As the number of protons in the nucleus of the atom increases, there is a greater force of attraction for the electrons in the shell and the distance between the electrons and the nucleus decreases. p+ 3p+ Ionization energy The energy required to remove an electron from a neutral atom. (The electron removed is the outermost and therefore the most easily removed electron and is always a valence electron unless the atom has a closed shell.) Possible to have 1 st ionization energy, 2 nd ionization energy etc. This refers to removing one electron (1 st ), then removing a 2 nd electron and so on. Ionization energy increases going across a row left to right Ionization energy decreases going down a group. Electronegativity Electronegativity is the tendency of an atom to attract electrons from a neighbouring atom. If an atom has high Electronegativity, it strongly attracts electrons from a neighbouring atom and may completely remove an electron from the neighbouring atom. This also means that atoms with high Electronegativity strongly attract their own valence electrons. As a result, these valence electrons are difficult to remove nd the atoms have high ionization energy. If an atom has low Electronegativity, it has little attraction to the electrons of a neighbouring atom and little tendency to remove an electron from a neighbour. Such an atom also has a small attraction to its own valence electrons. Therefore, these valence electrons are easier to remove and the atom has low ionization energy. The top right corner of the periodic table has the highest electronegativity (excluding the noble gases). Fluorine is the MOST electronegative element. Electronegativity increases as you go across a row from left to right. Electronegativity decreases as you go down a group. ***Use the Diagram below to help you remember the trends*** Simmer Down Inc. (Updated Dec 2012) 15

16 Electron Shell Recall that each period on the periodic table represents a different layer of electrons; that is, a different electron shell. The 1 st shell has 2 electrons and therefore the 1 st period has 2 elements. The 2 nd shell has 8 electrons and therefore the 2 nd period has 8 elements. The 3 rd shell has 8 electrons and therefore the 3 rd period has 8 elements. (We are going to ignore the transition metals, lanthanides and actinides) Open Shell: Shell containing less than its maximum number of electrons. Closed Shell: Shell contains its maximum number of electrons. Example: Look at the 2 nd shell (Lithium to Neon). The 2 nd shell holds 8 electrons (2s 2 2p 6 ). Neon is a noble gas and thus has a full outer shell. Therefore, Neon has a closed shell. Li to F all have less than 8 electrons in the outer shell and thus are considered open. Simmer Down Inc. (Updated Dec 2012) 16

17 Valence Electrons Recall that we said valence electrons were all the electrons in an atom except those in the core or in filled d- or f-subshells. We can rephrase this definition. Valence electrons are the electrons in open shells. Valence electrons are reactable electrons. The noble gases have no valence electrons and thus are not very reactive. Valence: number of electrons that are available for bonding or the number of unpaired electrons on the atom. Types of Chemical Bonding ***Valence is sometimes called combining capacity.*** The following must be remembered when examining trends across or down the periodic table. Ionic Bonds When going down a family in the periodic table, properties are affected by the increase size of the atoms and the increase distance between the nuclei and the valence electrons. When going across a period in the periodic table, properties are affected by the differing valence, nuclear charge and charge on the species. An ionic bond is formed by the attraction of a positive ion to a negative ion. An ionic bond is formed when an electron from one atom is transferred to another atom, so as to create one positive and one negative ion. How to predict when an ionic bond will form: Ionic bonds are formed when elements from opposite sides of the periodic table are combined. That is, when a metal and a non-metal are combined. Very important In general, when an atom forms an ion the atom loses or gains sufficient electrons to attain a closed shell. Covalent Bonding Covalent bond is a bond, which involves the equal sharing of electrons. Covalent bond is formed when two atoms having less than full shells of electrons are able to share one or more of their electrons with each other to attain full electrons. Octet Rule: Atoms in columns 14 to 17 of the periodic table tend to form covalent bonds so as to have eight electrons in their valence shells. Simmer Down Inc. (Updated Dec 2012) 17

18 How to predict if a covalent bond will form Covalent bonds are formed when a non-metal combines with a non-metal. Covalent bonds form when both atoms involved have relatively large electronegativities. In general, non-metals have large to very large electronegativities, attract each other s electron strongly, and will not let go of there own. Some covalent compounds have very high melting points. BN (boron nitride) melts at about 3000 C. SiC (Silicon carbide) melts at about 2700 C. These compounds form crystals with a network of covalent bonds. Makes each crystal one huge molecule held together by identical bonds. Not all covalent compounds have these high melting points. CH 4 melts at -182 C. O 2 melts at -218 C. Why do these compounds have such a low melting point? The answer lies in how each individual molecule is attracted to each other. In the previous examples (SiC and BN) there was covalent bonding between all the adjacent atoms. That is not the case with CH 4 and O 2. The covalent bonds that hold the individual atoms together are very strong, but the force that holds each molecule together is relatively weak. Writing Lewis Dot Structures Lewis Structures (electron dot structures) are used to help visualize the arrangement of bonds in molecules. The symbol is used to denote the nucleus and dots are used indicate the number of valence electrons. To write the Lewis Structure for an atom, write its chemical Li Ca B symbol surrounded by a number of dots which represent its valence electrons (electrons in outermost s and p orbitals). C N The Lewis Structure of an ionic compound is written by determining the charge for each ion and arranging the non-metal ions symmetrically around the metal ion. e.g. Draw the Lewis Structure for MgCl 2 Drawing Lewis Structures of covalent compounds that obey the octet rule (groups 14 to 17) follow a simple set of rules. Simmer Down Inc. (Updated Dec 2012) 18

19 Rules for drawing Lewis Diagrams are as follows 1. Count up the total number of valence electrons in the molecule. Adjust this number by SUBTRACTING one electron for every POSTIVE charge and ADDING an electron for each NEGATIVE charge on the molecule. 2. Determine which atoms are bonded together and put 2 electrons into each bond. 3. Use the remaining valence electrons to complete the octets of the atoms surrounding the central atom. Then place the remaining electrons, in pairs, on the central atom(s) (These non-bonding pairs of electrons are called lone pairs). 4. If the central atoms has less than an octet of electrons, have a neighbour share electrons with the deficient atom by putting an extra pair (or pairs) of electrons into the shared bond. 5. Last step; replace each pair of electrons engaged in a bond with a dash. Examples: 1. Draw the Lewis structure for NH 4 + How many valence electrons? N has 5, H s each have 1. Take one electron away for the positive charge and what do you have 8!!! Place the N in the middle and the H s around it. Don t forget the Octet rule 2. Draw the Lewis Structure for CHO 2-3. Draw the Lewis Structure for HOPO Simmer Down Inc. (Updated Dec 2012) 19