Chapter 12
By Ms.Cohane
Structure Determines Properties ●
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A cardinal principle principle of chemistry is that the macroscopic observed properties of a material are related to its microscopic structure, and vice-versa. The microscopic structure entails: ○ ○ ○ ○
the kinds of atoms the manner in which they are attached their relationship to other molecules the shape of the molecule
CHAPTER 12 OUTLINE ○
12.1 - Bond Types 12.2 - Electronegativity
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12.3 - Polarity
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12.4 - Electron Configurations 12.5 - Ionic Structures
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12.6 & 7 - Lewis Diagrams 12.8 - Bond Angles
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12.9 - Molecular Geometry
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12.10 - VSPER Theory
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GOALS ○
Explain how + why different types of bonds form
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Compare the types of bonds
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Describe how electrons are placed in an atom, and orbitals Create orbital diagrams, distinguish between different elements
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Create Lewis Structures for elements, and understand electrons further by applying Ch.11 skills
12.1 Types of Chemical Bonds
Bond: a Bond: a force that hold atoms together and makes them function as a unit
● Metal atoms lose electrons to become
positively charged ions – cations ● Nonmetals gain electrons and become
negatively charged ions – anions
“co” = sharing ● “valent” = valence electrons covalent bonds form between nonmetals ●
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The hydrogen atoms above are sharing two electrons. One from each original atom.
Think Pair Share ○ ○ ○ ○ ○ ○ ○ ○
Ion Bond Covalent Ionic Valence Electrons Orbitals Electron Diagram Orbital Diagram
“Polar” vs. “Nonpolar” Covalent
If e’ between atoms are shared equally equally,, the bond is classified as nonpolar covalent covalent
When atoms share e’ unequally unequally,, the bond is classified as polar as polar covalent covalent
11.11 Atomic Properties & Periodic Table
When a metal reacts with a non-metal, it transfers electrons from the metal to the non-metal
– The metal loses electrons and becomes a cation – The non-metal gains electrons, becoming an anion
11.11 Atomic Properties & Periodic Table Atomic radius – radius – the the size of an element’s atoms, usually the typical distance from the nucleus to the boundary of the surrounding cloud of electrons.
– Increases down each group Each step down a group an energy level is i s added
– Decrease from left to right across a period Electrons are added to the same energy level, but an increase in number of protons pulls electrons closer to nucleus
Atomic size of “parent” atom versus cation. A cation is smaller by an energy level than its parent atom
Atomic size of “parent” atom versus anion. An anion is slightly larger due to having more e’ than its parent atom
Electron Shielding – Shielding – electrons in the energy levels between the nucleus and the valence electrons are called "shielding" electrons because they "shield" the valence electrons from the force of attraction exerted by the positive charge in the nucleus.
Ionization Energy – Energy – the energy required to remove an electron from an atom • In main main group group elemen elements ts it it incre increase asess from from lef leftt to right right • Electrons are closer to nucleus • Non-metals gain electrons, metals lose electrons
12.2 Electronegativity
The power of an atom to attract electrons
Electronegativity values help us determine if a bond is nonpolar covalent, polar covalent or ionic.
What type of bond? Look Look at the difference in numbers between atoms!
• “polar” means unequal distribution of charge • “nonpolar” means equal sharing
Water - Polar Covalent
Practice ●
Classify the following bonds. F—H S—H H—H Na—Cl O—H
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Polar Covalent Polar Covalent Non-polar Covalent Ionic Polar Covalent
Rank the above bonds in order of increasing polarity
H-H, S-H, O-H, F-H, Na-Cl
12.3 Bond Polarity & Dipole Moments ●
Covalent bonding between unlike atoms results in unequal sharing of the electrons. We use the delta sign to show charge. δ+ H • F δ− •
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A molecule with an end end of positive charge and an end of negative charge is polar is polar & & has a dipole moment
Dipole Moment ○ ○
The magnitude of charges and the distance of separation between the charges. charges. In a water molecule, oxygen has a greater electronegativity than the hydrogen atoms, the electrons are not shared equally. So a charge distribution causes the molecule to behave like it has two centers centers of charge, one + and one _
center of positive charge
center of negative charge
Dipole Moment ●
The dipole moment affects the attractive forces between molecules, and therefore the physical properties too
● Water
molecules are strongly attracted to one another ❖ This is the reason for its high specific heat and high boiling point • Thank goodness water is not nonpolar – our oceans would be empty!
12.4 Stable Electron Configurations and Charges on Ions ●
Metals lose their valence electron(s) to form cations → The their new valence shell has the same electron configuration as the previous noble gas.
Nonmetals gain electron(s) to form anions → So their new valence shells have the same electron arrangement as the next noble gas.
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There have to be enough electrons from the metallic atoms to supply the needed electrons for the nonmetal atoms.
Properties of Ionic Compounds ●
All solids at room temperature (MP greater than 300°C)
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Liquid / molten state conducts electricity
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Brittle and hard
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Often soluble in water (KISS Guidelines)
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When dissolved in water the solution solution becomes an electrical conductor. All strong electrolytes because the ionic bonds don’t break!
12.5 Structure of Ionic Compounds
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Crystal lattice: geometric pattern determined by the size and charge of the ions
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Anions almost always larger than cation
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Anions generally considered “hard” “hard” spheres packed as efficiently as possible, with the cations occupying the “holes”
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Each cation is surrounded by as many anions as will fit
12.6 Lewis Structures ●
Bonding involves just the valence electrons of atoms.
Lewis Structure: a representation of a molecule that shows how the valence electrons are arranged among the atoms in the molecule. ● When bonding, the goal for each atom is to gain a noble gas configuration. configuration. ●
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The duet rule for rule for helium configuration. configuration. The octet rule for rule for other noble gases.
Lewis Dot Structures
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Use the symbol of the element to represent the nucleus and the kernel electrons Use dots around the symbol to represent valence electrons
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Put one electron on each side first, then pair up
In your notes, add in the next row, starting with Rb!
Lewis Structures Elements in the same group have the same Lewis structure, because they have the same number of valence electrons ● Cations have Lewis symbols without valence electrons. ● Anions have Lewis Lewis symbols with with 8 valence electrons. ●
Writing Lewis Structures Structures of Molecules Molecules Step 1 = 1 = count up all valence electrons in the molecule 2 = place the atoms in a symmetric pattern ● Step 2 = 3 = place single bonds (one pair of electrons) ● Step 3 = between the central atom and all outer atoms 4 = if any electrons are left over, satisfy s atisfy the octet ● Step 4 = rule for all atoms including the central atom ● Step 5 = check your work ○ Are all valence electrons counted for? ○ Do all atoms have an octet? ● Step 6 = if BOTH checks above don’t hold true, you need to create multiple bonds with the outer atoms ●
12.7 Lewis Structures of Molecules with Multiple Bonds Single covalent bond: bond: atoms share 2 electrons ● Double covalent bond: atoms share 4 electrons ● Triple covalent bond: atoms share 6 electrons ● Bond strength = triple > double > single ○ For bonds between same atoms, C≡N > C=N > C—N ○ Double is not 2x the strength of single, and triple is not 3x the strength of single ● Bond length = single > double > triple ○ For bonds between same atoms, C—N > C=N > C≡N ●
Problems with Lewis Structures ●
Some atoms do not tend to follow the octet rule.
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B and Be are often found octet-deficient Elements in the 3rd period or below often have expanded octets
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Some molecules have an odd number of electrons.
Draw Lewis Structures carbon tetrachloride – CCl 4 hydrogen sulfide – H2S nitrogen trihydride – NH 3 hydrogen cyanide – HCN carbon dioxide – CO 2 ozone – O3 2sulfate – SO4 nitrite – NO2
Resonance ●
Resonance Structures: when Structures: when there are multiple Lewis structures for a molecule that differ only in in the position of the electrons ○ Lone pairs & multiple bonds are in different places ○ The actual molecule is a combination of all forms
12.8 Molecular Structure Molecular Structure OR Geometric Structure: Structure: refers to the three dimensional arrangement of the the atoms in a molecule. Water is often called “bent” or “v-shaped”
To describe the structure more precisely, we often specify the bond angle – it is about 105° for water.
12.9 Valence Shell Electron Pair Repulsion (VSEPR Theory) A model for predicting the molecular structures of molecules formed from nonmetals ○ The structure around an atom is determined by minimizing repulsions between electron pairs. ○ Bonding pairs and lone pairs are positioned as far apart as possible from each other ○
Electron Pair Arrangement 180°
Linear
Trigonal planar
120°
109.5°
Tetrahedral
Molecular Structure ●
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Each bond counts as one area of electrons. Each lone pair counts as one area of electrons. ○ Even though lone pairs are not attached to other atoms, they do “occupy space” and repel other atoms more and affect angles. (They count in EPA)
When determining EPA, just count the number of quadrants that are filled with an atom or e’s
r of these below, they are tetrahedral (4) In each
Review… Electron Pair Arrangement & Resulting Molecular Structure (2,3,4 – pairs) Type CO2 BF3 O3 CH4 NH3 H2O
Electron Pair Arrangement
Ball & Stick Model
Molecular Structure
Bond angle
