Chapter 8: Basic Concepts of Chemical Bonding

Explain the chemical basis of physical properties of molecular compounds such as lower melting point, and inability to conduct electricity when dissolved in water. (8.3H)

(??) where is this even

Describe the essential features of ionic bonding including electron transfer to form ions and electrostatic attraction if ions to form a solid. (8.2C)

-usually result from metal+nonmetal bonds -involves complete transfer of electrons from one atom to another -attraction/bonding is primarily due to electrostatic forces between cations and ions with charges

Use average bond enthalpies to calculate ΔHrxn. (8.8CC)

1) Determine number of bonds and types of bonds. 2) Add the bond enthalpies for broken bonds and formed bonds. 3) Subtract formed bonds from broken bonds.

Determine bond order for bonds involved in resonance structures. (8.7X)

1) Find preferred structure. 2) Count number of bonds.

Depict the formation of ions using electron configurations and Lewis symbols and write the formula of the resulting ionic compound. (8.2D)

Anions lose electrons and cations gain electrons.

Rank similar covalent bonds according to their length and strength. (8.8AA)

As the number of bonds between atoms increases, the bond length grows shorter and the bond strength increases.

Determine average bond order when given a formula by drawing the Lewis structure and determining bond orders of individual bonds. (8.8Z)

Bond order: number of bonding pairs of electrons between atoms. Draw lewis structure and count the bonds between atoms. CN- bond order=3

Recognize the attractive and repulsive forces present in covalent bonds. (8.3I)

Bond strength is directly proportional to number of bonding electrons. Stronger bonds are shorter bonds.

Describe how Coulomb's law explains the periodic trend in lattice energy. (8.2F)

Coulomb's law: Eel=kQ1Q2/d As the distance increases, the electrostatic attraction decreases(lattice energy). As the charges increase, the attraction also increases.

Describe formation of covalent bond as resulting in a filled valence shell. (8.3G)

Covalent bond=sharing of electrons in bonds between atomic centers. Valence electrons around atoms form into either shared electron pairs as bonds or remain on one atom as lone pairs

Recognize that polarity of a molecule can be quantified using the idea of a dipole moment. (8.4O)

Dipole shows bond polarity with + end of arrow near S+ atom. Arrow head points to direction of electron pulling.

Describe how a reaction can be divided conceptually into bond-breaking and bond-forming steps. (8.8BB)

Each bond has a specific bond enthalpy. So when you add all the broken bond enthalpies and formed bond enthalpies and subtract them from each other, you are able to find the change in enthalpy of the reaction.

Understand that covalent and ionic bonding result in filling the outer valence level of each atom in a molecule. (8.1B)

Either with sharing electrons or a complete transfer of electrons.

Explain how bond polarity arises from differences in electronegativity of bonded atoms. (8.4L)

Electronegativity: the relative ability of a bonded atom to attract shared electrons to itself. The difference in electronegativity shows whether or not the electrons will be shared or transferred.

Use Lewis electron-dot symbols to depict main-group elements. (8.1A)

Elements in groups show valence electrons. Group 1A:1,2A:2,3A:3...

Calculate formal charges from Lewis structures. (8.5R)

Formal charge: a hypothetical charge on an atom in a structure assuming all atoms have the same electronegativity (so all bonding electrons are shared equally). FC=number of neutral atom's valence electrons-number of its non-bonding electrons-number of bonds

Explain how changes in bond strength account for the heat of reaction. (8.8DD)

Heat of reaction is calculated using bond enthalpy. Bond enthalpy: the measure of bond strength in a chemical bond. Changes in the bond strength will directly affect the heat of reaction.

Identify that polar bonds can contribute to overall polarity of molecules. (8.4N)

If the electronegativity is closer to 0, the bond is more non-polar. If the electronegativity is closer to 2, the bond is more polar.

Describe periodic trends in electronegativity. (8.4K)

Increasing up to the right

Use electronegativity differences to identify nonpolar covalent, polar covalent, and ionic bonds. (8.4M)

Ionic: EN>2.0 Polar Covalent: EN=0 Nonpolar Covalent: 0<EN<2.0

Define lattice energy and be able to arrange compounds in order of increasing lattice energy based on the charges and sizes of ions involved. (8.2E)

Lattice energy: the energy required to completely separate one mole of a solid ionic compound into its gaseous ion. smaller ionic radii=greater lattice energy(and vice versa)

Draw Lewis structures for compounds to show bonding between atoms and lone pairs in the molecule. (8.5P)

Lone pairs: pair of electrons not shared between atoms.

Draw Lewis structure when given molecular formula of hydrocarbons (including multiple bonds). (8.5Q)

Multiple bonds are needed.

Recognize that multiple bonds involve sharing of multiple pairs of electrons. (8.3J)

Multiple bonds=multiple shared electrons.

Apply the octet rule and its three major exceptions: (1) molecules with a central atom that has an electron deficiency, (2) an odd number of electrons, or (3) an expanded valence shell. (8.7V)

Octet rule: atoms tend to gain, lose or share electrons to acquire 8 valence electrons. Exceptions: 1) molecules with an odd number of electrons 2) molecules with an expanded valence shell 3) molecules with a central atom that has an electron deficiency (n=3)

Recognize molecules where resonance structures are needed to describe the bonding. (8.6T)

Resonance: the use of two or more lewis structures to represent a single molecule or ion. Molecules where the central atom is in n=3 period can break the octet rule.

Use formal charge to identify the dominant Lewis structure for a molecule or ion. (8.5S)

The preferred structure can be determined by the formal charges of each atom. The structure with the fewest and smallest atomic FC's, or the structure with any negative FC's on the most electronegative atom.

Draw the dominant resonance structures for a molecule. (8.6U)

Use formal charges and find preferred structure.

Explain the relationship between bond type (single, double, triple), bond strength, bond length, and bond energy. (8.8Y)

bond energy (bond enthalpy): the enthalpy required to break a specific bond in one mole of a gaseous substance. bond strength: single bond<double bond<triple bond bond length: triple bond<double bond<single bond Bond strength is directly proportional to number of bonding electrons.