Download Original

1
Chemistry Chapter 15 & 16 Study Guide
Table of Content
*Given Outline______________________________________________________Page 2-3
*Chapter 15 Textbook Summary Outline________________________________Page 4-5
*Chapter 16 Textbook Summary Outline________________________________Page 6-9
Study Guide Unit 5 Dr. Galante
Chemistry 10H
2
Chapters 15 & 16
1-Be able to characterize different types of solids: metallic (ex. Iron, etc.), ionic (ex.
NaCl, etc), molecular (ex. sugar, etc), network (ex. Diamond, etc). Also be able to
state the differences in their physical properties (Melting point, electrical
conductivity, hardness, etc.).
2-Know the difference between crystalline (orderly arrangement) and amorphous
(lacks order) solids.
3-Be able to describe the three types of cubic unit cells: simple cubic, body centered
cubic and face-centered cubic.
4-Be able to determine how many atoms/ions/molecules are present in each type of
cubic unit cell.
5-Be able to determine the coordination number of an ion from a picture of the unit
crystal.
6-Know what an alloy is and be able to identify what common material are alloys
(ex. Steel, bronze, stirling silver).
7-Know the difference between substitution alloys (atoms are replaced in the unit
cell) and interstitial alloys (atoms fit into spaces of the unit cell).
8-Be able to distinguish between valence and core electrons for any given element.
9-Know the key differences between covalent and ionic bonding. Be able to predict
from a formula whether a compound is ionic or covalent.
10-Be able to draw Lewis electron dot structures for a given compound (ionic or
molecular).
11-Be able to show the transfer of electrons to form ionic compounds from the Lewis
dot structures of atoms.
12-Understand the significance of the octet rule in assigning Lewis electron dot
structures.
13-Realize that certain elements (boron, etc.) do not always obey the octet rule.
14-Be able to calculate the formal charge for a given atom in a molecule.
FC=VE-LPE-1/2BE
15-Understand the significance of resonance structures.
16-Be able to predict the most likely resonance structure using formal charges.
17-Understand what delocalization is and how it applies to benzene (C6H6).
18-Know how to apply VSEPR theory to predict the geometry on a molecule.
19-Be able to show how atoms hybridize using electron configurations.
20-Be able to use the concept of orbital hybridization to describe the bonding form of atoms in
molecules. Know how to assign sp, sp2, sp3, sp3d, and sp3d2 hybridization to the appropriate
atoms, and know what geometry each of these hybridization forms will give the molecule :
(Octahedral, square pyramidal, square planar, trigonal bipyramidal, see-saw, T-shaped, trigonal
planar, linear, tetrahedral, trigonal pyramidal, bent, etc.)
21-Be able to draw all of these geometries using lines, wedges, and dashed lines.
Also be able to identify and show all bond angles between atoms.
22-Understand that lone pairs of electrons take up more room than bonding electrons
and the resulting bond angles can be distorted.
23-From a chemical structure, be able to determine the hybridization of each atom in
3
the structure.
24-Know the difference between sigma and pi bonds. (a single bond-1 sigma bond,
double bond-1 sigma and 1 pi bond, triple bond-1 sigma and 2 pi bonds)
25-Be able to recognize polar bonds in molecules using the concept of
electronegativity.
26-Be able to predict if a molecule is polar or non-polar.
27-Know the following intermolecular forces (dipole-dipole, dipole-induced dipole,
London dispersion forces, and hydrogen bonding) in liquids and solids. Know their
relative magnitudes - which one is strongest, which is weakest, etc.
28-Know how intermolecular forces effect: Solubility, melting point and boiling
point.
29-Be able to describe the major intermolecular forces acting upon a molecule or
atom or a group of molecules or atoms.
30-Know the requirements for hydrogen bonding - a hydrogen atom bonded to
oxygen, nitrogen or fluorine
31-Be able to describe the main features of molecular orbital theory. Realize that the
number of molecular orbitals always equals the number of atomic orbitals used in
the combining atoms. Know that some of the bonds will be bonding and others will
be antibonding.
32-Be able to draw molecular orbital diagrams for all diatomic molecules (Li2, O2,
NO, etc.) up to neon. Know how to handle a charged diatomic molecule (ex: O22-)
using molecular orbital theory. You will be provided with the correct molecular
orbital diagram if 2p electrons are used.
33-Be able to apply Hund’s Rule (when orbitals are of equal energy, electrons fill like
seats on the bus), Aufbau Principle (electrons fill in orbitals from lowest to highest
energy) and the Pauli Exclusion Principle (no more than 2 electrons per orbital
with opposite spins) when assigning electrons to molecular orbitals.
34-Be able to predict paramagnetic behavior (unpaired electrons in MOParamagnetic, all electrons paired in MO-diamagnetic).
35-Be able to calculate the bond order from molecular orbital diagrams.
BO=1/2(bonding electrons-antibonding electrons). Know what these predictions
mean for the molecule.
36-Be able to answer all questions from the quizzes, worksheets, classwork and
webassigns.
Textbook Outline
Chapter 15 Section 1: Electron Configuration in Bonding
Valence Electrons
*Definition: Electrons in the highest occupied energy level of an element’s atoms
*Determines chemical properties, determined (one way)through electron configuration or group#
*Exceptions to group # rule: noble gases with a full valence shell of 8 (Helium has 2)
4
*Valence electrons are the only electrons in chemical bonds (only valence in electron dot
structures)
Electron Configuration for Cations
* Lewis explained why atoms form certain ions and molecules through the octet rule, in the
formation of compounds atoms tend to achieve the electron configuration of a noble gas, a set
of 8 (except helium!) electrons in its highest energy level and general configuration ns2np6 .
* Metallic elements loose electrons making a complete octet in the next-lowest energy level
* Charges of cations can be explained by the loss of valence electrons by metal atoms (i.e. Na+ )
* Some pseudo noble-gas electron configurations: silver, copper, gold, cadmium, and mercury
Electron Configuration for Anions
*relatively full valence shell, gain electrons
*halide ions are ions produced when atoms of halogens gain electrons
Chapter 15 Section 2: Ionic Bonds
Formation of Ionic Compounds
*Ionic bonds are the forces of attraction between oppositely charged ions (cations&anions)
*Ionic compounds are compounds that consist of electrically neutral groups of ions joined by
electrostatic forces (do not have molecular formulas since not made up of molecules)
*Ex: NaCl (goal is for a stable octet)
*Formula units indicate the lowest whole-number ratio of cations to anions in any sample of an
ionic compound
Properties of Ionic Compounds
*At room temperature, most ionic compounds are crystalline solids arranged in repeating three
dimensional patterns such as NaCl
*Makes NaCl stable due to the large attractive forces of each sodium ion surrounded by 6
chloride ions and vice versa, reflected with ionic compounds high melting temperatures.
*Coordination number is the number of ions of opposite charge that surround the ion in a
crystal (Na+ =6, Cl-=6) (Cs+=8, Cl-=8) (TiO2=Ti4+ coordination #=6, O2-=3)
*In crystals, internal crystal structures differ, determined by a technique called x-ray diffraction
crystallography (x-rays pass through a crystal are recorded on film which shows the deflection
of x-rays in order to calculate the positions of ions in the crystal to define the structure)
*When melted or in aqueous solution, ionic compounds conduct an electric current
(electricity) due to the movement of ions when a voltage is applied (cations freely move to one
electrode and anions to the other; flow of electricity between the electrodes through an external
wire)
Chapter 15 Section 3: Bonding in Metals
Metallic Bonds and Metallic Properties
*Cations are surrounded by mobile valence electrons that drift from one part (of the metal) to
another
5
*Metallic bonds consist of the attraction of the free-floating valence electrons for the positively
charged metal ions (forces of the attraction that hold metals together)
*Good conductors of electrical current since electrons can flow freely
*Ductile (drawn into wires) and malleable due to the mobility of valence electrons
*Under pressure, metal cations easily slide past one another (unlike ionic crystals when struck
with a hammer)
Crystalline Structure of Metals
*Metals are also crystalline (even metals that contain one kind of atom!)
*Metals are arranged in compact and orderly patterns of
-body centered cubic (“8 neighbors”)
-face centered cubic (“12 neighbors”)
-hexagonal close-packed (“12 neighbors”) but rearranged
Alloys
*Alloys are mixtures composed of two or more elements (at least one which is metal)
*Generally prepared by melting a mixture of ingredients, then cooling
*Examples: brass=ally of copper+zinc; steel
*Important because their properties are superior to those of their component elements
*Forms from in two ways
-If the same size, replace each other in the crystal
-Different size, small atoms fit into the interstices (spaces) between larger atoms
(interstitial alloy aka steel)
Chapter 16 Section 1: The Nature of Covalent Bonding
Single Covalent Bonds
*Covalent bonding is like a “tug of war” where compounds such as HCl or H2O do not give up
or accept electrons as readily as NaCl
*Single covalent bond is a bond in which two atoms share a pair of electrons (such as the
diatomic molecule of hydrogen)
*Chemical formulas of ionic compounds describe formula units while chemical formulas of
covalent compounds describe molecules.
*correct formulas of molecular compounds reflect the actual number of atoms in each
molecule and their subscripts are not necessarily lowest whole-number ratios.
*Combinations of atoms of nonmetallic elements in Group 4A-7A are likely to form covalent
bonds. Lewis described this tendency in terms of the octet rule for covalent bonding
~“sharing of electrons occurs if the atoms involved acquire the electron configurations of
noble gases”
*Halogens form single covalent bonds in their diatomic molecules
6
*When carbon forms bonds with other atoms, it usually forms 4 bonds with the promotion of the
2s electrons to the vacant 2p orbital (the stability compensates for the small energy cost of
electron promotion)
Coordinate Covalent Bonds
*A coordinate covalent bond is a covalent bond in which one atom contributes both bonding
electrons (like any other covalent bond but different source)
*Atoms in polyatomic ions are covalently bonded, contain covalent and coordinate covalent
bonds (the negative charge shows number of electrons in addition to valence electrons present)
Bond Dissociation Energies
*When hydrogen atoms combine into hydrogen molecules, a large quantity of heat is released
*The release of energy suggests that the product is more stable than the reactants
*Bond dissociation energy is the total energy required to break the bond between two
covalently bonded atoms
*Compounds with high dissociation, therefore are pretty unreactive, include compounds with
only C-C and C-H single covalent bonds (like methane).
Resonance
*Previously, chemists imagined that electron pairs alternate and readily flip back and forth, or
resonate (using a double headed arrow to indicate)
*Double covalent bonds are usually shorter than single bonds so it was believed that bond
lengths (such as in ozone) were unequal
*However, it was later discovered that electron pairs do not actually resonate back and forth,
but form a hybrid, or mixture, of the extremes represented by the resonance forms
*Resonance structures are structures that occur when it is possible to write two or more valid
electron dot formulas that have the same number of electron pairs for a molecule or ion.
Exceptions to the Octet Rule
*For some molecules or ions, it is impossible to write structures that satisfy the octet rule when
the total number of valence electrons is an odd number
*Example: NO2 has 17 valence electrons
-Therefore has two resonance structures with 6 bonding electrons, and 11 lone electrons
-It is impossible to write a Lewis dot structure NO2 that satisfies the octet rule for all
atoms, yet the nitrogen dioxide molecule does exist as a stable molecule
*Diamagnetic are substances in which all of the electrons are pairs (the magnetic effects of
paired electrons essentially cancel, weakly repelled by an external magnetic field)
*Paramagnetic are substances that contain one or more unpaired electrons (show a strong
attraction to an external magnetic field)
*Example of resonance structure: O2. Hybrid of two structures (one with a single covalent bond,
10 loan electrons and one with a double covalent bond and 8 loan electrons)
*Compounds that expand the octet to include 10 or 12 electrons: PCl5 (opposed to PCl3 which
follows the octet rule) and SF5
Chapter 16 Section 2: Bonding Theories
Molecular Orbitals
7
*Model for covalent bonding has to do with the assumption that atomic orbitals of individual
atoms are unchanged in bonded atoms (quantum mechanical model of bonding)
*Parallels: atomic orbitals belong to a particular atom, molecular orbital to a particular whole
molecule, two electrons to fill both an atomic and a molecular orbital
*molecular orbitals or orbitals that apply to the entire molecule form when two atoms combine
and their atomic orbitals overlap
-requires # of molecular orbitals=# of overlapping atomic orbitals
-when 2 atomic orbitals overlap, 2 molecular orbitals are created
1) bonding orbital, a molecular orbital with an energy that is lower than that of the atomic
orbitals from which it formed (electrons seek lower energy first-so it gets filled up first)
2) antibonding orbital, an energy that is higher than that of the the atomic orbitals from
which it formed
*Example: hydrogen molecule (H2)
-the 1s atomic orbitals of the 2 hydrogens overlap
-2 electrons (1 from each separate hydrogen) are available for bonding
-the energy of the electrons goes into the bonding molecular orbital and therefore is
lower than the energy of the electrons in the atomic orbitals of separate hydrogen atoms
-result is a stable covalent bond between the hydrogens
-notice: antibonding orbital is EMPTY
-high probability of finding electrons between the nuclei of the bonded atoms
-a sigma bond, when two atomic orbitals combine to form a molecular orbital that is
symmetrical along the axis connecting two atomic nuclei, is formed
*covalent bonding results from an imbalance between the attractions and repulsions of the
nuclei and electrons involved (tug of war)
-attractions between the hydrogen nuclei and the electrons are stronger than the
repulsions of nuclei and other nuclei as well as electrons and other electrons (creating a stable
diatomic molecule of H2!)
*However, in antibonding orbitals combining two atoms leads to electron pairs in higher energy.
These are not found between in the nuclei, and the balance between attractions and repulsions
favors the repulsions (such as the He2 molecule).
*Atomic p orbitals as well as s orbitals overlap (such as the F2 molecule-end-to-end overlap)
*A pi bond is formed by side-by-side overlap of p orbitals; bonding electrons are most likely to
be found in sausage-shaped regions above and below the bond axis of the bonded atoms
*Pi bonds are weaker than sigma bonds
VSEPR Theory
*Valence-shell electron pair repulsion theory states that because electron pairs repel, molecular
shape adjusts so the valence electron pairs are as far apart as possible
*VSEPR theory tells you molecular shapes
*Example: methane molecule has 4 bonding pairs, farthest at 109.5º(tetrahedral)
*Unshared loan pairs, such as NH3, are held closed to the nitrogen and strongly repels the
bonding electrons pushing them closer (making the angle 107º)
-also applies to H2O, with two loan pairs, compressing the angles to about 105º
-in contract, CO2 has no loan pairs so it is linear and 180º
Hybrid Orbitals
8
*To describe the types of bonds formed is hybridization, several atomic orbitals mix to form
the same total number of equivalent hybrid orbitals
*Useful in describing double covalent bonds, such as ethene
-sp2 hybridized from the combo of one 2s and two 2p atomic orbitals of carbon
-120º
*Page 457-459 have plenty of examples!
Chapter 16 Section 3: Polar Bonds and Molecules
Bond Polarity (ALL BASED ON ELECTRONEGATIVITY!!!!)
*Covalent bonds differ in terms of how bonded atoms share electrons
-depends on the kind and number of atoms joined together
*nonpolar covalent bonds=the bonding electrons are shared equally when the atoms in the
bond pull equally (such as H2, O2, N2)
*polar covalent bond (polar bond) forms when a covalent bond joins two atoms of different
elements and the bonding electrons are shared unequally
-more electronegative atom has stronger electron attraction=slightly negative charge
*Ex: HCl is polar, since chlorine is more electronegative and has the stronger electron attraction
the “polarity arrow” would be pointing in a set direction (with chlorine having a slightly -charge)
*Ex: H2O is polar since the highly electronegative oxygen pulls the bonding electrons away from
hydrogen in a set direction
*If the electronegativity difference is greater than 2.0, high probability that electrons will be
completely pulled away by on of the atoms and an ionic bond will form
Polar Molecule
*In a polar molecule, one end of the molecule is slightly negative and the other end is slightly
positive
*A molecule that has two poles is called a dipolar molecule or dipole (such as HCl)
*when dipoles are placed in an electric field, they orient with respect to the positive and negative
plates
*A molecule such as carbon dioxide can have polar bonds, but because it is linear the bond
polarities cancel out.
Attraction Between Molecules
*molecules are attracted to each other by intermolecular attractions
-these attractions are weaker than either an ionic or a covalent bond
-responsible for determining the state of a molecular compound
1) Van der Waals forces are the weakest attractions, consist of:
a) dispersion force: weakest of all molecular interactions, cause by the motion of electrons,
increases as # of electrons increase
a) Ex: Attraction between halogen diatomic molecules is through dispersion forces
b) Fluorine and chlorine have few electrons, week dispersion forces=gases at STP
b) dipole interactions: occurs when polar molecules are attracted to one another
a) Electrostatic attractions involved occur between oppositely charged regions of dipolar
molecules
b) slightly negative region of a polar molecule is attracted to the slightly positive of anot.
9
c) similar but much weaker to ionic bonds
2) Hydrogen bond
*In water, dipole interactions result in weak attraction of water molecules for one another
*Hydrogen bonds are attractive forces in which a hydrogen covalently bonded to a very
electronegative atom is also weakly bonded to an unshared electron pair of another
electronegative atom. STRONGEST OF INTERMOLECULAR FORCES
-this atom can be in the same molecule or in a nearby molecule
-hydrogen bonding always involves HYDROGEN
~why? because it is the only chemically reactive element with valence electrons
that are not shielded from the nucleus by a layer of underlying electrons
-very polar covalent bonds form why hydrogen bonds with oxygen, nitrogen, or fluorine
-sharing of a nonbonding electron pair on a nearby electronegative atom compensate for
the deficiency of electrons
-hydrogen bond has about 5% the strength of a covalent bond
**Chart on page 465 that compares characteristics of ionic and covalent compounds**