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Chemistry 11 > Chapter 6 - Claire > Flashcards

Chapter 6 - Claire Flashcards

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1
Q

What Δ EN values create an ionic bond?

A

1.7 to 3.3

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2
Q

What Δ EN values create a covalent bond?

A

0 to 1.7

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3
Q

Where is atomic radius the largest?

A

Bottom left of the periodic table

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4
Q

Are cations smaller or larger than the neutral atom?

A

Smaller

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5
Q

Are anions smaller or larger than the neutral atom?

A

Larger

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6
Q

Def. the net positive charge experienced by an electron in a multi electron atom, representing less than the full nuclear charge as inner electrons repel outer electrons, offsetting some of the nucleus’s inward pull(I.e. shielding effect)

A

Effective nuclear charge

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7
Q

Def. the repulsive action of inner core electrons in outer valence electrons, effectively reducing the nucleus’s inward pull, causing the atom to increase in size

A

Shielding effect

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8
Q

Coulomb’s law states that the _______ of force depends directly on the amount of ____________

A

Strength, electrostatic charge

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9
Q

Which two elements have no shielding effect?

A

Hydrogen and helium

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10
Q

What are the four types of charge?

A
  1. Proton-proton repulsion
  2. Neutron-proton attraction (residual strong force)
  3. Electron-electron repulsion
  4. Nucleus-electron attraction
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11
Q

Def. an expression of the size of an atom, representing the typical distance from the center of its nucleus to the boundary of its electron cloud

A

Atomic radius

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12
Q

Def. an estimate of the size of an atom, based on half the distance between covalently bonded atoms

A

Covalent radius

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13
Q

Def. an estimate of an the size of a metal atom based on half the distance between the nuclei of adjacent atoms in a metal lattice

A

Metallic radius

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14
Q

What affects an electron cloud/ an atoms size?

A
  1. Number of occupied shells

2. Amount and type of charge

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15
Q

Def. predictable and regular patterns of cyclical change in elemental properties throughout the periodic table, moving either vertically or horizontally

A

Periodic trends

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16
Q

Def. the amount of energy required to remove the most loosely held electron from a gaseous atom to form a cation, or from a cation to increase its positive charge

A

Ionization energy

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17
Q

Where are electrons held the most loosely?

A

Bottom right of the periodic table

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18
Q

Def. a chemical property that describes an atom’s (or a group of atoms’) ability to attract electrons towards itself

A

Electronegativity

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19
Q

Def. the change in energy when an atom gains an electron to form an anion, or when an anion gains another electron, increasing its negative charge

A

Electron affinity

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20
Q

The higher the electronegativity, the higher the ____________

A

Electron affinity

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21
Q

def. A numerical value representing the difference in the electronegativities of two atoms

A

ΔEN

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22
Q

def. A covalent bond between non-metal atoms with the same electronegativities, where
electron sharing is equal between them

A

Nonpolar covalent bond

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23
Q

Are electrons ever transferred to a different atom?

A

No, they only effectively transfer because there is always some sort of electronegativity

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24
Q

def. A covalent bond between two non-metal atoms in which electron sharing is sufficiently
unequal that the bond displays significant and meaningful ionic nature

A

Polar covalent bond

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25
Q

Where is the electron sharing is equal or nearly equal between the atoms and the
bond has predominantly covalent character with little to no ionic character?

A

Nonpolar bonds

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26
Q

Where is the electron sharing is disproportionately uneven enough that the bond has
predominantly ionic character, with very little covalent character?

A

ionic bonds

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27
Q

def. The non-integer charge values at the ends of a polar covalent bond, due to the uneven
distribution of electron density within the bond

A

Partial ionic charges (δ+, δ–)

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28
Q

Where is the electron sharing is uneven enough that

the covalent bond has some ionic nature – more than a nonpolar bond, but less than an ionic?

A

polar bonds

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29
Q

what Δ EN values form a polar bond?

A

0.4-1.7

30
Q

what Δ EN values form a non-polar bond?

A

0-0.4

31
Q

def. A simplified diagram that communicates molecular structure using:
▪ An element’s symbol to represent its nucleus and core electrons
▪ Dots to represent any nonbonding electrons in the valence shells
▪ Lines to represent the pairs of bonding electrons in chemical bonds

A

Lewis structure

32
Q

What are Lewis structures of individual atoms called?

A

electron dot structures

33
Q

when orbitals hybridize they become ________

A

equivalent

34
Q

def. describes an atom’s combining power, often in terms of how many chloride anions it can ionically bond with

A

valency

35
Q

for metals, valency represents how many electrons it will _______

A

lose

36
Q

for non-metals, valency represents how many electrons it will _______

A

gain

37
Q

Apart from how many electrons a non-metal will gain, what does valency represent for non-metals?

A

how many covalent bonds it will form

38
Q

the lewis structure of _______ compounds show the component ions and the ratio between them

A

ionic

39
Q

for the lewis structures of ionic compounds, you must surround the ions with _______ brackets and ___________

A

square, superscript to represent their ionic charges

40
Q

ionic compounds exist in what way?

A

ionic crystal lattices

41
Q

covalent compounds exist in what way?

A

as separate and distinct molecules

42
Q

Valence electrons can be either _______ or _______

A

bonding or non-bonding

43
Q

def. any pair of valence electrons that are non-bonding, as atoms don’t share them with other atoms

A

lone pair

44
Q

lone pairs may _______ or _______ because of bond formation

A

form or disperse

45
Q

hydrogen obeys the ____ rule in molecules

A

duet

46
Q

def. an observational rule in chemistry that says the valence shells of the first-period elements hydrogen and helium can hold a maximum of two electrons, never more

A

duet rule

47
Q

def. An observational rule of thumb in chemistry that says main-group elements from
periods 2 to 7 seek to have eight electrons in their valence shell

A

octet rule

48
Q

What are the exceptions to the octet rule?

A

Boron, Beryllium, Aluminium, Hydrogen and Helium

49
Q

What periods can exceed the octet rule?

A

Periods 3-7

50
Q

In a Lewis structure, for binary compounds which element is usually the central atom?

A

the one with the lower electronegativity

51
Q

def. Covalent bonding between two atoms who share more than two electrons

A

Multiple bond

52
Q

def. A configuration of elements from periods 3 through 7, in which their valence shells
expand past eight electrons as the central atom, as they can incorporate d orbitals

A

expanded octet

53
Q

when a molecule’s central atom takes on more than eight electrons, what do we call the molecule?

A

Hypervalent molecule

54
Q

def. Different, but equivalent Lewis structures of the same molecule, showing all the
possible double bond locations, connected by a double-headed arrow

A

resonance structures

55
Q

def. Electrons in a molecule, ion or solid metal that aren’t associated with a single atom or
covalent bond, allowing them to be co-valent to more than two atoms

A

Delocalized electrons

56
Q

When atoms share electrons unevenly enough that the the bond takes on ionic character, the atoms develop _________

A

partial charges

57
Q

the _____ electronegative atom gains electron density and develops a partial negative charge δ–

A

more

58
Q

the ______ electronegative atom loses the electron density, and develops a partial positive charge δ+

A

less

59
Q

def. describes the separation of opposite charge

A

electric dipole

60
Q

def. A vector quantity that expresses a bond’s polarity, showing the separation of opposing
charges (an electric dipole), and the directionality of that separation (a moment)

A

bond dipole moment

61
Q

def. o A molecule that has an overall, or net dipole moment, because of its bond polarities and
its molecular geometry

A

polar molecule

62
Q

def. A vector quantity expressing a molecule’s overall polarity, being the net dipole moment
resulting from the vector addition of the molecule’s individual bond dipole moments

A

molecular dipole moment

63
Q

when can a molecule be polar?

A

A molecule can only be polar if it contains polar bonds and

is asymmetrical, in either shape, composition or both

64
Q

def. Attractive forces that exist within a molecule or compound, holding the atoms together

A

Intramolecular forces

65
Q

def. Forces, mainly attractive, that exist between molecules and neighbouring particles, be
they other molecules, atoms or ions

A

Intermolecular forces

66
Q

def. The attractive force that exists between the partial positive charge of one polar
molecule and the partial negative charge of another

A

dipole-dipole force

67
Q

dipole-dipole forces exist between polar molecules that have ______ molecular dipole moments

A

permanent

68
Q 
def. An electrostatic attraction that forms between the partially positive hydrogen in a highly
polarized bond (usually H–N, H–O and H–F) and the negative lone pair on a nearby atom
▪ Stronger than a typical dipole-dipole interaction because it has some covalentbond
character of electron sharing
A

Hydrogen bond

69
Q

In polar molecules, quantum mechanical uncertainty affects the
________ of the molecular dipole moment

A

strength

70
Q

In nonpolar molecules, quantum mechanical uncertainty affects the ________ of a
molecular dipole

A

existence

71
Q

def. A weakly attractive force that forms when nonpolar molecules gain temporary dipole
moments

A

London dispersion force

Chemistry 11 (6 decks)

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