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A1 Flashcards

atomic structure, isotopes, electrons, ionic bonding

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1
Q

isotope

A

atoms of the same element
same number of protons
different number of neutrons

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2
Q

element

A

substance made up of only 1 type of atom
cant be broken down into another substance

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3
Q

atom

A

smallest particle of an element that can exist

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4
Q

Relative masses of:
- proton
- neutron
- electron

A

1
1
1/1836

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5
Q

relative charges of:
- proton
- neutron
- electron

A

+1
0
-1

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6
Q

atomic number

A

the number of protons

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7
Q

mass number

A

sum of protons + neutrons

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8
Q

why do isotopes have the same chemical properties

A

they have the same number of electrons

same arrangement of electrons
neutrons have no effect on chemical properties

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9
Q

anion

A

negative ion
formed when atom GAINS electrons

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10
Q

cation

A

positive ion
formed when atom LOSES electrons

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11
Q

relative atomic mass

A

weighted mean mass
of an atom of an element
compared to 1/12th of the mass of an atom of carbon 12

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12
Q

relative isotopic mass

A

weighted mean mass
of an atom of an isotope
compared to 1/12th of the mass of an atom of carbon 12

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13
Q

binary compound

A

contains only 2 elements

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14
Q

polyatomic ions

A

ions which have more than one type of atom

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15
Q

how many electrons do these shells hold?
- 1st shell
- 2nd shell
- 3rd shell
- 4th shell

A

2
8
18
32

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16
Q

how many electrons do these subshells hold?
- s
- p
- d
- f

A

2
6
10
14

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17
Q

why do we write 4s before 3d

A

because 4s is in a lower energy level than 3d
we write starting with lowest energy level

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18
Q

orbital

A

region within an atom
around the nucleus
which holds up to 2 electrons with OPPOSITE SPINS

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19
Q

what do the arrows represent in electrons in boxes

A

spin

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20
Q

give the shape of and s orbital and a p orbital

A

spherical
dumbell

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21
Q

what are the 2 exceptions to electron configurations

A

chromium
copper

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21
Q

ionic bonding

A

electrostatic attraction
between oppositely charged ions

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21
Q

describe structure and physical properties of ionic bonding

A
  • giant ionic lattice
  • high melting/boiling points
  • soluble in POLAR solvents
  • solid at room temperature
  • brittle
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22
Q

what 2 factors affect ionic bonding strength

A

ionic charges
ionic radii

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23
Q

explain how ionic charges and radii size affect ionic bonding

A

stronger ionic charges= stronger attraction= stronger ionic bond

smaller radii= stronger attraction= stronger ionic bond

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24
Q

how does an ionic compound dissolve in water

A

ionic lattice breaks down
the ions separate

water molecules surround ions

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25
Q

why don’t ionic compounds conduct electricity when solid

A

ions are in fixed position
and cannot move through the structure

26
Q

why do ionic compounds conduct electricity when molten/ aqueous

A

ions are mobile are free to move through structure and carry charge

27
Q

define molten

A

liquified by heat

28
Q

define aqueous

A

dissolved in water

29
Q

why do ionic compounds have high boiling/melting points

A
  • giant ionic lattice
  • strong ionic bonds
  • these break during melting
  • require lots of energy to overcome
30
Q

covalent bonding

A

shared pair of electrons

electrostatic attraction between shared pair of electrons and positive nuclei

31
Q

covalent bond is localised- what does this mean?

A

the attraction is soley between the shared pair of electrons and nuclei of the bonded atoms

32
Q

difference between ionic and covalent bonding

A

in ionic, electrons are transferred, in covalent electrons are shared

33
Q

lone pair

A

a pair of electrons that is not shared

34
Q

multiple bond

A

2 atoms share more than 1 pair of electrons

34
Q

double bond

A

attraction is between 2 shared pairs of electrons and nuclei of bonding atoms

35
Q

triple bond

A

attraction is between 3 shared pairs of electrons and nuclei of bonding atoms

36
Q

dative bond

A

both electrons in the shared pair come from 1 of the bonding atoms

37
Q

properties of simple molecular substances

A
  • low melting/boiling points (weak IMFs require little energy to overcome)

-dissolve in NON-POLAR solvents (IMFs form between non-polar solvent and simple molecule)

-don’t conduct electricity (no free electrons/ ions)

-gases

38
Q

properties of giant covalent structures

A
  • high melting/boiling points (strong covalent bonds require high energy to overcome)

-solids

-insoluble in polar and non-polar solvents (covalent bonds too strong)

  • don’t conduct electricity (except graphite)
39
Q

why does graphite conduct electricity

A

only uses 3 of its 4 bonds

has a free, delocalised electron to move through the structure and carry charge

40
Q

electron pair repulsion theory

A

electrons have negative charges so electron pairs repel each other as far as possible

lone pairs repel more than bonding pairs as they’re more electron dense (decrease bond angle by 2.5)

41
Q

octet rule

A

not all elements achieve a full outer shell of 8
- may be too few/much electrons

elements beyond group 15 and period 3 can expand their octet

42
Q

why are non linear shapes and trigonal bipyramids POLAR

A

-dipoles are not symmetrically arranged
- effects don’t cancel out
-there is an overall dipole

43
Q

state bond angles within
-2 regions
-3 regions
-4 regions

A

2 regions= 180 degrees
3 regions= 120 degrees
4 regions= 109.5 degrees

44
Q

electronegativity

A

the ability of a bonded atom to attract the shared pair of electrons in a covalent bond

45
Q

how is electronegativity measured

A

pauling scale

46
Q

how does number of protons and atomic radius affect electronegativity

A

more protons= increase electronegativity (electrostatic attraction towards electrons is greater)

small distance between outer shell and nucleus= increase electronegativity (stronger attraction)

47
Q

why are noble gases not electronegative

A

they have a full outer shell of electrons
don’t attract electrons

48
Q

what makes a bond non-polar? (pure covalent bond)

A

both atoms have same electronegativities
electron pair is shared evenly

49
Q

what makes a bond polar?

A

one atom is more electronegative
it attracts the shared pair of electrons
the shared pair is unevenly shared

50
Q

why are C–H bonds non-polar?

A

carbon and hydrogen have very similiar electronegativites

51
Q

in MgI there is polarisation- explain this term (3)

A
  • Mg and I have different electronegativities
  • electron pair is shared unevenly
  • there is an uneven distribution of charge
52
Q

state the 3 types of intermolecular forces

A

-induced dipole-dipole
-permanent dipole-dipole
-hydrogen bonding

53
Q

describe the origin of induced dipoles (london forces)

A

electrons in the electron cloud are constantly moving
at any moment, there is an instantaneous dipole
more electrons on one side
this instantaneous dipole induces a temporary dipole in a neighbouring molecule

54
Q

describe how number of electrons affects strength of london forces

A

more electrons means a larger instantaneous dipole and larger induced dipole
stronger attractive forces between molecules

55
Q

what does a hydrogen bond act between

A

lone pair on an electronegative atom and a hydrogen atom

56
Q

state anomolous properties of water

A

ice is less dense than water so it floats
relatively high Mp/Bp
high surface tension

57
Q

explain why ice floats on water

A

ice is less dense than water
because hydrogen bonds hold the water molecules apart in an open lattice

each water molecule can form 4 hydrogen bonds as there are 2 lone pairs on the O and 2 H atoms

58
Q

explain why water has a relatively high melting/boiling point

A

relatively strong hydrogen bonds which require large amount of energy to overcome

59
Q

why does boiling point increase going down the noble gases

A

there are more electrons
so stronger london forces
more energy required to overcome

60
Q

state the formula and charge of:

-nitrite
-hydrogencarbonate
-manganate(VI)
-silver

A

NO2-
HCO3-
MnO4-
Ag+

61
Q

state the formula and charge of:

-sulfite
-dichromate
-phosphate
-sulfide

A

SO3 2-
Cr2O7 2-
PO4 3-
S 2-

62
Q

metallic bonding

A

electrostatic attraction between cations and delocalised electrons

63
Q

describe the properties of metals

A

high Mp/Bp (strong metallic bonds)

conduct electricity (delocalised electrons)

insoluble (strong metallic bonds) solid at room temp

malleable (layer of cations slide over each other)

ductile

giant metallic lattice

chem exam questions (11 decks)

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