Ch. 5&6 Unit Test

Front 1

In the three molecules, O2 , HCl and F2 , what atom would have a partial negative charge?

Back 1

chlorine

Front 2

The electrostatic attraction between positively charged nuclei and negatively charged electrons permits two atoms to be held together by a(n)

Back 2

chemical bond.

Front 3

In metals, the valence electrons

Back 3

are shared by all of the atoms.

Front 4

How many electrons must be shown in the Lewis structure of the hydroxide ion,OH- ?

Back 4

8

Front 5

As atoms bond with each other, they

Back 5

decrease their potential energy, thus creating more-stable arrangements of matter.

Front 6

In drawing a Lewis structure, each nonmetal atom except hydrogen should be sorrounded by

Back 6

8 electrons.

Front 7

Dipole-dipole forces are considered the most important forces in polar substances because the London dispersion forces present in polar substances

Back 7

are usually much weaker, than the dipole-dipole forces.

Front 8

Compared with molecular bonds, the strength of intermolecular forces is

Back 8

weaker.

Front 9

According to VSEPR theory, the shape of an AB3 molecule is

Back 9

trigonal-planar.

Front 10

A neutral group of atoms held together by covalent bonds is a

Back 10

molecule.

Front 11

The lattice energy is a measure of the

Back 11

strength of an ionic bond.

Front 12

Compare with ionic compounds, molecular compounds

Back 12

have lower melting points.

Front 13

The properties of both ionic and molecular compounds are related to the

Back 13

strengths of attraction between the particles in the compounds.

Front 14

The concept that electrostatic repulsion between electron pairs surrounding an atom casues these pairs to be separated as far as possible is the foundation of

Back 14

VSEPR theory.

Front 15

The electron configuration of nitrogen is 1s²2s²2p³. How many more electrons does nitrogen need to satisfy the octet rule?

Back 15

3

Front 16

The reason the boiling point of water (H O) is higher than the boiling point of hydrogen sulfide (H S) is partially explained by

Back 16

hydrogen bonding.

Front 17

The chemical formula for an ionic compound represents the

Back 17

ratio of the combined ions present in a sample.

Front 18

The electrons involved in the formation of a chemical bond are called

Back 18

valence electrons.

Front 19

Use VSEPR theory to predict the shape of the carbon tetraiodide molecule, CI .

Back 19

tetrahedral

Front 20

Metals are malleable because the metallic bonding

Back 20

allows one plane of ions to slide past another.

Front 21

The equal but opposite charges present in the two regions of a polar molecule create a(n)

Back 21

dipole.

Front 22

Shifting the layers of an ionic crystal causes the crystal to

Back 22

shatter.

Front 23

The strength of London dispersion forces between molecules depends on

Back 23

both the number of electrons in the molecule and the mass of the molecule.

Front 24

Four hybrid sp³ orbitals are formed from

Back 24

one s orbital and three p orbitals.

Front 25

The B­-F bond in BF (electronegativity for B is 2.0; electronegativity for F is 4.0) is

Back 25

ionic.

Front 26

If two covalently bonded atoms are identical, the bond is

Back 26

nonpolar covalent.

Front 27

To draw a Lewis structure, one must know the

Back 27

number of valence electrons in each atom.

Front 28

Chemists once theorized that a molecule that contains a single bond and a double bond split its time existing as one of these two structures. This effect became known as

Back 28

resonance.