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SCH4U Exam Study Notes Essay

Electron configuration: notation that shows number and arrangement of electrons in its orbitals Infinite number of electron configurations because infinite values of n For each atom, all but one of these represents the atom in an excited state Atom’s chemical property mainly associated with its ground state electron configuration Unless otherwise stated, assume electron configuration is in its ground state Ex: 1s2 (superscript represents 2 electrons in s orbital)

Aufbau principle: process of building up the ground state electronic structure for atoms (in order of atomic #)

Writing electron Configurations:
Provide info about first 2 quantum numbers (n and l)
Ex: boron in ground state:

Orbital diagram: boxes are used for each orbital Empty box represents an orbital with no electrons, box with single upward or downward pointing arrow represents 1 electron, box with 2 oppositely pointing arrows represents a filled orbital (2 electrons)

Writing Electron Configurations for Periods 1 and 2:
Energy of each orbital increases form left to right
Fifth electron can go into any orbital (-1, 0, +1) in 2p sublevel because all have same energy but customary to place electron in first box

Electron Configurations and Orbital Diagrams for Period 3:
Same process as before but can use a simplified notation called a condensed electron configuration Since chemical reactivity depends mainly on atom’s valence electrons Electron configuration of noble gas of previous period is put in square brackets using atomic symbol only then continue with configuration of next energy level being filled Ex: Phosphorus: [Ne] 3s23p5

Electron Configurations and Orbital Diagrams for Period 4:
Electrons do not occupy 3d orbital until 4s orbital is filled because 4s orbitals have lower energy levels than the 3d orbital (fill orbitals in order of increasing energy levels) Chromium and copper have an unexpected deviation (only one electron in 4s) Because greatest stability results from a configuration where 4s is half filled (stable ground state)

Electron Configuration and the Periodic Table:
Arranging elements by electron configurations
S and p block elements called main group elements (wide range of physical and chemical properties) D block elements called transition elements (mark transition from p orbital to d orbital) F block elements called inner transition metals (mark transition from d orbital to f orbital)

Patterns Involving Group Numbers and Period Numbers Elements in a group have similar chemical properties because they have similar outer electron configurations ( same # of valence electrons) 3 patterns

1. For main group, last numeral of group number is same as number of valence e- Ex: strontium in group 15 and 5 valence electrons Exception is helium (2 valence electrons and in group 18 but energy level is complete) 2. N value of highest occupied energy level is period number 3. Square of n value (n2) equals total # of orbitals in that energy level Maximum number of electrons in any principle energy level is 2n2 (max 2 e- in each orbital) Matches number of elements in the period (does not work for period 3 because only 8 elements)

Summarizing Characteristics of s, p, d, and f block Elements: S block elements span 2 groups (groups 1 and 2)because s orbital can hold 2 electrons P block span groups 13 to 18 and p orbital can hold 6 electrons D block spans 10 groups because the d orbital can hold 10 electrons In general: ns, (n-1)d (chromium and copper are exceptions)

F block spans 14 groups because the f orbital can hole up to 14 electrons In general: ns, (n-2)F, (n-1)d, np

Electron Configurations, Atomic Properties, and Periodic Trends: Electron configurations help to determine atomic and chemical properties of elements Atomic radius, ionization energy, metallic character, and electron affinity are periodic Periodic trends in atomic radius

Can be determined by measuring distance between nuclei of bonded atoms Decreases across a period and increases down a group
2 factors
1. Increases with increasing n (as n increases, higher probability of finding electrons farther from nucleus) 2. Effective nuclear charge: net force attraction between electrons and nucleus (Zeff) Increases when there are more electrons and they are more strongly attracted to the nucleus and the atomic radius decreases Valence electrons experience a smaller Zeff because inner electron shield them from the nucleuses attractive force For main group elements (s and p block)

n governs increasing atomic radius down a group and Zeff decreases in valence electrons because of additional shielding Zeff governs decreasing atomic radius across a period (n is constant, Zeff increases) Transition elements (d and f block) do not display same trend Electrons added to inner energy levels (d orbital) instead of outer energy levels and Zeff remains fairly constant Periodic trends in ionization energy (electron is considered to be completely removed from its atom) Ionization energy: energy needed to completely remove one electron from a ground state gaseous atom Energy needed to remove an e- in order to overcome force of attraction from the nucleus Multi electron atoms have more than one ionization energy (ionization energies increase).

Atoms with low IE1 (first ionization energy) tend to form cations during reactions (group 1 has low IE1) Elements with high IE1 tend to form anions (- charged ions) (exception are noble gases) IE generally decreases down a group (atom radius increases, valence e- distance from nucleus increases, decrease in attraction of e- to nucleus ) Generally increases across a period (as Zeff increases, attraction of valence e- to nucleus increases) Reactivity of metals increases down a group and decreases across a period (in general) Group 12 are exceptions because have a configuration similar to a noble gas Periodic trends in electron affinity

Electron affinity: change in energy that occurs when an e- is added to a gaseous atom More than one electron affinity (EA1 results in formation of anion with 1- charge) High e- affinity = gains an electron easily (ex. Fluorine)

High negative numbers mean high electron affinity
Trends more irregular because other factors affect it
Elements in group 17 and 16 have high ionization energies and high electron affinities (attract e- strongly, form – ions) Elements in group 1 and 2 have low ionization energies and electron affinities (attract e- poorly, form + ions) Elements in group 18 have high ionization energies but low electron affinities (do not give up or gain e-) Pg 163-189 Chemical Bonding and Lewis Structures

What Bonding Is
Bonding is a way to gain stability
Systems with lower energy are more stable than systems of higher energy Bonded atoms have less energy than unbounded
Using Lewis Structures to Represent Bonding Atoms

Interactions of valence (highest energy level) electrons represented in Lewis structures Dots used for lone pairs
Filled clockwise
Bonding and Properties
Ionic Bonding
∂E ≥ 1.7 large difference results in ionic bonds
Usually metal (sdf) and non-metal (p)
Arranged in crystal lattice structure; repeating units
High melting point, very strong bond
Insulators as solid
Do not conduct heat, good conductors in solution
Hard but brittle crystals

Lattice Energy and Ionic Bonding
Lattice Energy: amount of energy released when gaseous ions of its elements bond to make crystal In KJ/mol
Same amount of energy must be added to break crystal back into ions Process that gives off energy increases stability
Why high melting point, higher lattice energy means higher melting point Metallic Bonding
Metals do not have good hold on electrons (delocalised electrons) So small forces of attraction occur between metal atoms, since electrons are attracted to other nucleus Responsible for ductile and malleable properties of metals, why pure metals can conduct electricity Covalent Bonding

Lowest Energy bonds, low melting and boiling point
Nonconductive even in solution (polar is better conduction)
Coordinate covalent bond (one electron contributes two bonding electrons) Like in nh4 where n contributes 2 electrons (dative bond)
Can be polar or non polar
Balance between the forces of attraction and repulsion between two atoms Optimum space in between
Overlapping of orbitals causes formation of new orbital with less energy than original orbitals Since electrons occupy lowest energy level, new orbital more stable for them Lewis Structures
Rules of Representing Bonding
X-X Single bond (2e-)
X=X Double Bond (4e-)
X(3) X Triple Bond (6e-)
.. Lone pair of e-
[A]+ [X}-

Steps
1. Count total number of e-
2. Pick an A (central atom, usually fewest, odd duck)
3. Place PA (Peripheral Atoms)
4. Place single bonds between CA and PA (keep running e- tally) 5. Fill PA Valence shells
6. Call Kaiwen Song’s house and order Chinese food (and a Richard) 7. Take care of CA last (may make double bond)
8. Tally at 0 when finished
..
O—Cl—O –
|
ClO3-
Resonance structure, more than one possible structure
Bonds moving within a molecule, very stable
Benzene has a resonance structure

Valence Bond Theory
Linus Pauling looked at how the quantum model of the atom relates to bonding

Sigma Bond
H-H bonds by 1s orbitals overlapping to create combined 1s-orbital, called a Sigma bond (2e-, between an s and an spd or f, or two p orbitals) (probability regions) H2O: oxygen has 2p4 while 2 hydrogens have 1s1. Combine the 2p and 1s to make 2 sigma bonds (1 per hydrogen) Sigma bond is the orbital overlap, quantum version of Lewis single bond Hybridisation

Be, Mg, Ca, Sr, Ba have full 2s shell, so why aren’t they like noble gases Hybridisation (cross between two things)
In Be for example, the 2s electrons moves to the empty 2p orbital, the s and p shells then lose and gain energy to equalise in energy, hybrid sp orbital created, then it bonds sp2 1 e- form s 2e- from p Only changes when it is going to bond due to energy change

When e- leaves a full s to an empty p the energy of both equalises, this is hydridisation for sp (dp, df fp all kinds of hybrid orbitals) Carbon about to bond: 1s: 2sp3:

Pi Bond
Quantum version of a Lewis double bond (4e- involved) Pi bonds are made from the partial hybridisation(some of th p e- stay in the p energy level instead of going down to the s) between the orbitals of two bond atoms The two hybridised p orbitals bond to make a pi bond and the two hybrid sp that’s how it is a double bond

Triple Bonds: Partial Hybridisation (sp and pp)

When an sp hybridises with two p e- 2 pi bonds and one sigma, for 3 bonds total pi and triple bonds must occur between overlapping ps (for now) VSEPR (Vesper) Theory Valence Shell Electron (outside electrons) Pair (lone pairs and bonded pairs 2e- are negatively charged) Repulsion (The negative charges repel) Theory (A darn tootin’ swell idea by the Englishman Glutsby) Law of electric charges (likes repel, opposites attract, neutral and charged attract) The charged ends (bonded electrons and lone pairs) distance from each as far as they can due to repulsion Steric number= number of bonds made by central atom+ lone pairs When you have more than one central atoms, you pick the two on the end and break it up E number of lone pair

AXmEn where A is central atom, X is other atom, n is number of lone pairs

Chemistry Unit 2 Review: Electrochemistry
Electrochemistry:
The branch of chemistry dealing with electric energy during electron transfer What are Oxidation states? Every atom and ion has its oxidation state, showing if it has gained or lost electrons in a reaction Elements in their free states (like diatomic HOBRFINCL gases) have an oxidation state of 0 A positive oxidation state means that the atom has become a cation, while negative oxidation state means it had become an anion The oxidation state of oxygen in a peroxide is -1 instead of -2 Functional groups and radicals have total oxidation states corresponding to their net charge Compounds have total oxidation state of 0

To find the oxidation state of an atom in a compound, consider the oxidation state of another atom in the compound, and find the number that adds that to make 0. Remember to consider the number of atoms and charges.

Ex. Find the oxidation state of Chlorine in AgCl2
The Ag has a total oxidation of 2, and -2+2=o
So the oxidation state of the chlorine must be -1, since -1(2)=-2

For functional groups, remember that the total oxidation state is the net charge

Ex: Find the oxidation state of Manganese in MnO4-
Oxygen’s total oxidation state is -2(4) or -8
The total oxidation must be -1, so Manganese must be in an oxidation state of 7 What is Oxidation and Reduction?
Oxidation and reduction refer to the change in oxidation states during reactions Loss and gain of electrons
Reactions in which oxidation and reduction occur are called redox reaction Remember LEO GER (Losing electrons, oxidising, gaining electrons reduction) Methods of Balancing Redox Reactions
Half Reaction Method
The half reaction method works by splitting a redox reaction into an oxidation and a reduction The goal is to have the same amount of electrons on the opposite sides Remember, when there is an oxidation, there is also a reduction and vice versa Using the Half Reaction Method under neutral pH conditions


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