CH1010_Fall 2015_Dr. Kreider-­‐Mueller CH1010 Exam #2 Study Guide
For reference see “Chemistry: An Atoms-­‐focused Approach” by Gilbert, Kirss, and Foster Chapter 4: Chemical Bonding
•
Definitions to know:
o Molecule: unit of matter that results when 2 or more atoms are joined by covalent bonds.
o Ionic Bond: a bond resulting from the electrostatic attraction of a cation for an anion.
o Nonpolar Covalent Bond: a bond characterized by an even distribution of charge; electrons
in the bonds are shared equally by the two atoms.
o Polar Covalent Bond: a bond resulting from unequal sharing of bonding pairs of electrons
between two atoms.
o Polyatomic ions: charged group of more than one kind of atom joined together by covalent
bonds.
o Electrostatic Potential: the energy a charged particle has due to its position relative to another
charged particle.
§ Directly proportional to the product of the charges of the particles
§ Inversely proportional to the distance between the particles
o Lattice Energy: the energy released when 1 mole of an ionic compound forms from its free
ions in the gas phase.
o Bond Energy: the energy needed to break 1 mole of a particular covalent bond in a molecule
or polyatomic ion in the gas phase.
o Octet Rule: atoms of main group elements make bonds by gaining, losing, or sharing
electrons to achieve a valence shell containing 8 electrons, or four electron pairs.
o Bond Length: distance between the nuclear centers of two atoms joined together in a bond.
o Bonding Pair: a pair of electrons shared between two atoms.
o Lone Pair: a pair of electrons that is not shared between two atoms.
o Electronegativity: a relative measure of the ability of an atom in a bond to attract electrons to
itself when bonded to another atom.
o Resonance: characteristic of electron distributions when two or more equivalent Lewis
structures can be drawn for one compound.
o Resonance Structure: one of two or more Lewis structures with the same arrangement of
atoms but different arrangements of bonding pairs of electrons
o Resonance Stabilization: the stability of a molecular structure due to delocalization of its
electrons.
o Formal Charge: value calculated for an atom in a molecule or polyatomic ion by determining
the difference between the # of valence electrons in the free atom & the sum of lone-pair
electrons plus half of the electrons in the atom’s bonding pairs.
o Bond Order: the number of bonds between atoms: 1 for a single bond, 2 for a double bond, 3
for a triple bond.
•
Be able to Distinguish between the three different types of bonds
* Energy is released when bonds are formed*
* Energy is absorbed when bonds are broken*
1) Nonpolar Covalent Bond: two atoms evenly share 2 electrons
o Length and Strength of covalent bonds:
§ Double bond is shorter and stronger than a single bond
§ Triple bond is shorter and stronger than a double bond
1 of 12 CH1010_Fall 2015_Dr. Kreider-­‐Mueller § The internuclear distance that corresponds to a potential energy minimum (point
“c” on the graph below) is the bond length between two atoms.
2) Polar Covalent Bond: two atoms share 2 electrons unevenly (unequal sharing). The
electrons are more attracted to the more electronegative atom.
3) Ionic Bond: complete transfer of one or more electrons from one atom to another occurs
resulting in the formation of 2 charged particles (ions)
o Ions of opposite charge attract one another
§ Metals tend to form cations
§ We use a Roman numeral in parentheses to indicate # of charges on a cation
• Ex: in Fe(II)Cl2 the iron has a 2+ charge, Fe2+
• Ex: in Fe(III)Cl3 the iron has a 3+ charge, Fe3+
§ Nonmetals tend to form anions
§ Common polyatomic ions you should know:
•
NH4+, ammonium
NO3−, nitrate
ClO4−, perchlorate
CH3CO2−, acetate
OH−, hydroxide
PO43−, phosphate
CN−, cyanide
MnO4−, permanganate
SO32−, sulfite
NO2−, nitrite
CO32−, carbonate
SO42−, sulfate
Be able to distinguish between ionic & molecular compounds in molecular level representations
Ionic
Covalent
2 of 12 CH1010_Fall 2015_Dr. Kreider-­‐Mueller • Be able to account for the differences in physical properties of ionic & covalent compounds.
For ex, see table 4.1 from General Chemistry: Atom’s First by McMurry & Fay
•
Trends to know (and be able to explain…think about Zeff):
o Lattice Energy
§ Depends on the charge of the ions and the size (radius) of the ions.
• The smaller the radius, the higher the lattice energy
• The larger the charge on the ions, the higher the lattice energy
§ The higher the lattice energy, the higher the melting point of the ionic solid
§ Be able to predict the magnitudes of lattice energy based on analogous compounds,
ionic charge, & ionic radius
• For example, see the table below:
o Electronegativity
§ Look at Figures 4.8 & 4.9 in your textbook
§ Be able to predict the ionic/covalent character of bonds based on electronegativity
differences
• Bonds btw atoms w/ similar electronegativities are usually nonpolar covalent
• Bonds btw atoms whose electronegativities differ by more than 2 units are
largely ionic
• Bonds between atoms whose electronegativities differ by more than 0.4 units
but less than 2 units are usually polar covalent
§ Fluorine is the most electronegative (value of 4.0) atom
§ Oxygen is the second most electronegative (value of 3.5) atom
3 of 12 CH1010_Fall 2015_Dr. Kreider-­‐Mueller • Bonding in molecules can be represented by using only valence electrons in electron dot (Lewis)
diagrams. Be able to draw electron dot structures:
o Section 4.3 of your book gives some steps for drawing electron dot structures
o Figure out how many valence electrons you need to account for in your structure
o Connect the atoms appropriately
o Remember the following:
§ Hydrogen does not follow the octet rule (only surrounded by 2 electrons)
§ C, O, N, and F follow the octet rule (boron often does not)
§ If an element is in period 3 or below, then you can expand the octet if necessary
§ Carbon likes to form 4 bonds
§ Nitrogen typically likes to form 3 bonds
§ Oxygen typically likes to form 2 bonds
§ Halogens typically like to form 1 bond
o Consider if you need any multiple bonds
o If there is more than one possible structure, you need to be able to determine which
structure is best based on the octet rule and formal charges
Formal Charge = (# of valence e−) – ½(# of bonding e−) – (# of nonbonding e−)
§ The more formal charges there are on the atoms in your structure, the less the
structure contributes to the overall structure of the molecule (more formal charge =
worse structure)
o Consider if you can draw resonance structures for your compound (ex. Ozone, O3)
§ If you can draw different resonance structures, then the electrons are delocalized
within the compound
§ Delocalization reduces electrons’ potential energy making the molecule more stable
§ Be able to calculate bond orders for resonance hybrids
Chapter 5: Bonding Theories, Explaining Molecular Geometry
*Bonding can be described using different models…each is useful & each has drawbacks.*
• Definitions to know:
o Bond Angle: the angle (in degrees) defined by lines joining the centers of two atoms to a
third atom to which they are chemically bonded.
o Electronic Geometry: the 3D arrangement of bonding pairs & lone pairs of electrons
about a central atom.
o Molecular Geometry: the 3D arrangement of the atoms in a molecule.
o Bond Dipole: separation of electrical charge created when atoms with different
electronegativities form a covalent bond.
o Hybridization: in valence bond theory, the mixing of atomic orbitals to generate new sets
of orbitals that then are available to form covalent bonds with other atoms.
o Hybrid Atomic Orbital: in valence bond theory, one of a set of equivalent orbitals about
an atom created when specific atomic orbitals are mixed.
o Molecular Orbital: a region of characteristic shape and energy where electrons in a
molecule are located.
o Bonding Orbital: term in MO theory describing regions of increased electron density
between nuclear centers that serve to hold atoms together in molecules.
o Antibonding Orbital: term in MO theory describing regions of electron density in a
molecule that destabilize the molecule because they do not increase the electron density
between nuclear centers.
4 of 12 CH1010_Fall 2015_Dr. Kreider-­‐Mueller • Understand the Valence-Shell Electron-Pair Repulsion (VSEPR) model
o Used to predict the shape of a molecule
o Electrons in bonds and in lone pairs are thought of as occupying “electron domains”
(regions of electron density) that repel one another and stay as far apart as possible,
causing the molecules to assume specific shapes
o When applying the VSEPR model:
§ Draw the electron-dot structure
§ Be able to predict the geometric arrangement of electron domains around each atom
by assuming that its electron domains are oriented in space as far away from one
another as possible
•
Know electron and molecular geometries for molecules with 2, 3, 4, 5, or 6 electron domains
o Be able to predict bond angles based on the shape of a molecule or polyatomic ion
§ For trigonal planar and tetrahedral electron geometries, as the # of lone pair
electrons increases, the bond angles become more compressed (bond angle
decreases)
§ Ex: CH4, NH3, and H2O all have a tetrahedral electronic geometry (4 electron
domains). CH4 has no lone pair electrons, and the H-C-H bond angle is about
109.5˚. NH3 has one set of lone pair electrons on the N, and the H-N-H bond angle
is about 107˚. H2O has 2 sets of lone pair electrons on the O, and the H-O-H bond
angle is about 104.5˚.
5 of 12 CH1010_Fall 2015_Dr. Kreider-­‐Mueller o Be able to identify simple polar molecules based on their molecular shape & bond polarity.
§ Polar Bonds: bond between 2 atoms that have a large difference in electronegativity § Polar Molecule: A molecule that has a significant net dipole moment § A molecule that contains polar bonds may or may not be a polar molecule!! • You have to consider the geometry of a molecule in order to determine if it has a net dipole moment • If a molecule contains polar bonds, but the dipole moments for those bonds point equally in opposite directions, the dipole moments cancel and there is no net dipole moment on the molecule (ex. CO2) § Examples of polar molecules: H2O, CH3OH, CH2Cl2, NH3 § Examples of nonpolar molecules: CO2, SiH4, CH4, CCl4, CH3CH3 • Valence Bond Theory: provides a readily visualized orbital picture of how electron pairs are
shared in a covalent bond
o Section 5.4 of your book
o Covalent bonds are formed by overlap of atomic orbitals, each of which contains 1 electron
of opposite spin. The 2 overlapping lobes must be of the same phase
o Each of the bonded atoms maintains its own atomic orbitals, but the electron pair in the
overlapping orbitals is shared by both atoms
o The greater the amount of orbital overlap, the stronger the bond
o Hybridization
§ What does the concept of hybridization help us explain?
§ Know about the shape (dumbbell shape with one large lobe & one small lobe) and
orientation of sp, sp2, and sp3 hybrid orbitals
§ Be able to identify the hybridization on a given atom in a molecule.
• Ex: The carbon atom in CH4 is sp3 hybridized
• Ex: Each of the carbon atoms in C2H4 are sp2 hybridized
• Ex: each of the carbon atoms in C2H2 are sp hybridized
§ You can determine the hybridization of an atom based on the number of electron
domains
• See Table 5.3 in your book.
• If you have sp3 hybridization, you mixed 1 s atomic orbital with 3 p atomic
orbitals so you have no unhybridized p atomic orbitals
• If you have sp2 hybridization, you mixed 1 s atomic orbital with 2 p atomic
orbitals so you have one unhybridized p atomic orbital
• If you have sp hybridization, you mixed 1 s atomic orbital with 1 p atomic
orbitals so you have 2 unhybridized p atomic orbitals
§ Be able to identify the orbitals that overlap to give a specific bond in a molecule
• A σ bond involves “head-on” overlap
• A π bond involves “side-on” overlap
• A single bond consists of 1 σ bond
• A double bond consists of 1 σ and 1 π bond
• A triple bond consists of 1 σ and 2 π bonds
6 of 12 CH1010_Fall 2015_Dr. Kreider-­‐Mueller • Know a few basic things about Molecular Orbital Theory
o Atomic Orbital: a wave function whose square gives the probability of finding an electron
within a given region of space in an atom
o Molecular Orbital: A wave function whose square gives the probability of finding an
electron within a given region of space in a molecule
o Molecular orbitals are to molecules what atomic orbitals are to atoms
o A molecular orbital describes a region of space in a molecule where electrons are most
likely to be found, and it has a specific size, shape, and energy level
o MO’s are formed by combining atomic orbitals on different atoms. The # of molecular
orbitals formed is the same as the # of atomic orbitals combined
o MO’s that are lower in energy than the starting atomic orbitals are bonding, and MO’s that
are higher in energy than the starting atomic orbitals are antibonding.
o Electrons occupy MO’s beginning with the MO of lowest energy. A maximum of 2
electrons can occupy each orbital, and their spins are paired
o Bond order can be calculated: B.O. = [(# bonding e−)-(# of antibonding e−)]/2
o Be familiar with the a basic MO diagram: ex for H2, H2+, He2
• Combine the concepts of Valence Bond Theory and Molecular Orbital Theory to explain
resonance hybrids & the delocalization of electrons.
Chapter 6: Intermolecular Forces, Attractions Between Particles (Sections 6.1-­‐6.3 only) •
Definitions to know:
o Intramolecular Forces: any force that holds together the atoms making up a molecule or compound. They contain all types of chemical bonds. o Intermolecular Forces: forces of attraction or repulsion between neighboring particles o Dipole Moment: a quantitative expression of the polarity of a molecule. o Temporary Dipole: an intermolecular force between nonpolar molecules caused by the
presence of temporary dipoles within the molecules. o Permanent Dipole: permanent separation of electrical charge in a molecule due to unequal distribution of bonding and/or lone pairs of electrons. o Van der Waals Force: Any interaction between neutral atoms and molecules including hydrogen bonds, other dipole-­‐dipole interactions, and London dispersion forces. (Does not apply to interactions involving ions) o London Dispersion Force: an intermolecular force between nonpolar molecules caused by the presence of temporary dipoles in the molecules. o Ion-Dipole Interaction: an attractive force between an ion and a molecule that has a permanent dipole. o Dipole-Dipole Interaction: an attraction between regions of polar molecules that have partial charges of opposite sign. o Hydrogen Bond: the strongest dipole-­‐diploe interaction, which occurs between a hydrogen atom bonded to a N, O, or F atom and another N, O, or F atom. 7 of 12 CH1010_Fall 2015_Dr. Kreider-­‐Mueller o Polarizability: the relative ease with which the electron cloud in a molecule, ion, or atom can be distorted, inducing a temporary dipole. o Sphere of Hydration: the cluster of water molecules surrounding an ion in an aqueous solution. o Alcohol: An organic compound whose molecular structure includes a hydroxyl group bonded to a carbon atom that is not bonded to any other functional group(s). o Solvent: the component of a solution that is present in the largest # of molecules. o Solute: any component in a solution other than the solvent. A solution may contain one or more solutes. o Solubility: the maximum quantity of a substance that can dissolve in a given volume of solution. o Miscible: Liquids that are mutually soluble in any proportion. o Hydrophobic: describes a “water-­‐fearing” or repulsive interaction between a solute & water that diminished water solubility. o Hydrophilic: describes a “water-­‐loving” or attractive interaction between a solute and water that promotes water solubility. • Be able to do the following: o Identify polar molecules based on their molecular shape and bond polarity o Distinguish between intermolecular & intramolecular forces § Be able to identify the 4 kinds of intermolecular forces (IMF’s) • Ion-­‐dipole o Act between ions & molecules o Result of electrical interactions between an ion & the partial charges on a polar molecule • Dipole-­‐Dipole (A Van der Waals force) o Experienced by neutral but polar molecules o Result of electrical interactions among dipoles of neighboring molecules o Strength of force increases as the polarity (net dipole moment) of the molecule increases • London Dispersion (A Van der Waals force) o Experienced by all atoms & molecules, regardless of structure o Due to motion of electrons (instantaneous dipole) o Magnitude of force depends on the polarizability of the molecule o The larger/heavier the molecule, the more polarizable it is, & the stronger the dispersion forces o The more spread-­‐out (longer chain) the molecule is, the more surface area there is available for dispersion forces o Generally a weak force 8 of 12 CH1010_Fall 2015_Dr. Kreider-­‐Mueller • Hydrogen Bonding (A Van der Waals force) o Attractive force between a hydrogen atom bonded to a very electronegative atom (O, N, F) & an electron-­‐rich region elsewhere, either in the same molecule or in a different molecule §
o Be able to explain why hydrogen bonds are significantly stronger than dipole-­‐dipole forces o Be able to discuss the significance of hydrogen bonding in water The table below presents a comparison between these forces (Table from General Chemistry: Atom’s First by McMurry & Fay) o IMF’s are responsible for physical properties of materials o IMF’s can be predicted from molecular structure, shape, & bond polarity o Physical properties can be predicted from molecular structure § As the strength of IMF’s increases, boiling point increases • Ex: the larger the atomic radius, the greater the polarizability of the electron clouds. The greater the polarizability, the more likely the atoms/molecules are to form temporary dipoles. The stronger the London Dispersion Forces, the higher the boiling point. (Table 6.1 & 6.2) 9 of 12 CH1010_Fall 2015_Dr. Kreider-­‐Mueller o Ex: For a series of hydrocarbons, the longer the hydrocarbon chain, the greater the polarizability of the electron clouds. The greater the polarizability, the more likely the atoms/molecules are to form temporary dipoles. The stronger the London Dispersion Forces, the higher the boiling point. See the figure below. o As the strength of the London Dispersion forces increases, solubility decreases § For example: be able to recognize solubility trends for alcohols • As the hydrocarbon chain length increases, the solubility decreases Alcohol Solubility in Water (g/100g) CH3OH Infinitely soluble CH3CH2OH Infinitely soluble CH3CH2CH2OH Infinitely soluble CH3CH2CH2CH2OH 9 CH3CH2CH2CH2CH2OH 2.7 CH3CH2CH2CH2CH2CH2OH 0.6 CH3CH2CH2CH2CH2CH2CH2OH 0.18 CH3CH2CH2CH2CH2CH2CH2CH2OH 0.054 CH3CH2CH2CH2CH2CH2CH2CH2CH2CH2OH Insoluble in water 10 of 12 CH1010_Fall 2015_Dr. Kreider-­‐Mueller Chapter 7: Stoichiometry, Mass Relationships & Chemical Reactions * Your best method of studying this chapter is to practice problems! Look at the mock exam, old
exam, and any sample problems we have gone over in class.*
•
Definitions to know:
o Stoichiometry: the mole ratios among the reactants and products in a chemical reaction
o Chemical Equation: a description of the identities & proportions of reactants and products
o Reactants: substances consumed during a chemical reaction
o Products: substances formed during a chemical reaction
o Molecular mass: the mass in amu of one molecule of a molecular compound
o Molar mass: the mass of 1 mole of a substance
§ A convenient measure for comparing quantities of substances because the number
of atoms/molecules in a typical sample is huge!
o Formula mass: the mass in amu of one formula unit of an ionic compound
o Law of conservation of mass: the principle that the sum of the masses of the reactants in a
chemical reaction is equal to the sum of the masses of the products
•
Be able to do the following:
o Balance chemical equations. Your textbook outlines the following steps that can be used to
balance chemical equations:
1) Write an express using correct chemical formulas that includes one mole of the known
reactants & products
§ Include symbols indicating physical states
2) Check whether the expression is already balanced. If it is, then stop.
3) If the expression is not balanced, choose an element that appears in only one reactant &
product to balance first. Insert the appropriate coefficient(s) to balance this element
4) Choose the element that appears in the next fewest reactants & products & balance it.
5) Repeat the process for additional elements if necessary
§ Mass and number of atoms are conserved in a chemical reaction (Law of conservation
of mass)
§ Be able to interpret chemical reactions conceptually:
• Ex: The balanced chemical equation for the reaction below is:
2A2 + B2 →2A2B
o Calculate grams of a compound if you are given moles of a compound
§ Moles can be used to convert between molecular level particles & macroscopic masses
of compounds
o Calculate moles of a compound if you are given grams of a compounds
o Use the chemical formula of a compound to calculate the number of atoms in a sample of
a given compound
§ Be able to relate mass or moles of the compound to the mass or moles of the elements
that comprise the compound
11 of 12 CH1010_Fall 2015_Dr. Kreider-­‐Mueller o Perform stoichiometric calculations given a balanced chemical equation
§ Mol/mol ratios allow us to convert between different compounds
§ Be able to perform mass to mole conversions
§ Be able to perform mole to mass conversions
§ Such calculations are summarized in the table below:
12 of 12