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physical chem > bonding > Flashcards

bonding Flashcards

(94 cards)

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1
Q

substances

A

made of single repeating unit at atomic level

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2
Q

mixture

A

made of several repeating units at atomic level

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3
Q

2 properties of particles

A

always have kinetic energy
always hold onto each other

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4
Q

3 properties of metals

A

malleable
ductile
conductive

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5
Q

electron sea model

A

electrons in a metal can move freely around in the structure

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6
Q

what keeps electrons together in electron sea model

A

positive ions form structure and balance charge of delocalized electrons

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7
Q

whats a lattice

A

regular repeating structure of ions

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8
Q

what holds lattice together

A

electrostatic force between cation and electron

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9
Q

stronger attraction in cations and electrons

A

stronger metallic bonding

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10
Q

why are metals malleable

A

lattice shape changed but cations can slide over each other keeping shape and metallic bonding

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11
Q

what happens when metal turns to liquid

A

regular lattice structure broken and cations can slide over each other

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12
Q

properties of liquid metal

A

not malleable or ductile
shiny and conductive

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13
Q

bonding in a liquid metal

A

does have delocalized electrons
weaker metallic bonding

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14
Q

higher melting point of metals

A

stronger metallic bonding

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15
Q

properties of gas

A

no metallic bonding and higher boiling point

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16
Q

do metals in liquid state have lattice

A

no lattice of cations

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17
Q

what happens when a metal boils

A

overall attractive forces overcome repulsive forces so no balance of forces

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18
Q

why do metals have higher melting and boiling points

A

because there is a strong attraction between cations and delocalized electrons

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19
Q

properties of metal gases

A

only have metal atoms

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20
Q

trend of metallic bonding

A

increases across period

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21
Q

why do metallic bonding increase across period

A

ions in lattice become more charges and number of electrons increases so higher electrostatic force between cation and electrons

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22
Q

why cant metals delocalize electrons all the time

A

only delocalize outer shells electrons

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23
Q

why can metals only delocalize outer shell electrons

A

only largest orbital overlap
small orbitals dont overlap so cant delocalize

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24
Q

trend of metallic bonding down group

A

decreases

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25
why does metallic bonding decrease down group
size increases so more distance between cations nucleus and electrons so weaker forces
26
2 features of cations that effect metallic bonding
charge and size
27
state and explain general properties of metals exam question
melting point across period 3 eg charge of cation increases so stronger attraction between cations and electrons stronger attraction means stronger metallic bonding
28
compare melting point of 2 metals
eg mg ions more charges than potassium ions so stronger attraction of cations and electrons stronger bonding so more energy to break forces
29
explain general melting point down group
size increases so more distance between cation and electrons so less energy required
30
what are alloys
metals with 2 or more elements
31
properties of all salts
high melting and boiling points brittle soluble only conductive soluble
32
whats the perfect ionic model
all salts are made of positive and negative ions arranged in 3D lattice and ions perfect spheres
33
what keeps lattice together in ionic bonding
negative and positive electrostatic force/ attraction
34
what is ionic bonding
attraction between all cations and anions in lattice
35
charge of silver
+
36
charge of zinc
2+
37
charge of iron
2+ or 3+
38
charge of copper
1+ or 2+
39
charge of lead
2+ or 4+
40
group 7 6 5 anions
7= - 6=2- 5=3-
41
how to name anions
add ide to element name
42
hydrogencarbonate ion
HCO3-
43
sulfate ion
SO42-
44
phosphate ions
PO43-
45
maganate ions
MnO42-
46
hydroxide ions
OH-
47
how do cations and anions combine
1 to 1 ratio to keep compound neutral
48
how to present ionic compounds
use empirical formula
49
how ionic compounds are formed in terms of electrons
how many electrons transferred from one element to other
50
what happens when ionic compound melts
lattice breaks down and flow over each other weaker ionic bonding
51
liquid ionic compound into gas
seperates into individual compounds no ionic bonding
52
properties of solid ionic compounds
cant carry charge dont conduct electricity
53
properties of liquid ionic compound
can carry charge can conduct electricity
54
properties of gas ionic compiund
cant carry charge or conductive
55
why are ionic bonds so strong
attraction of negative and positive repulsion between both same
56
why do ionic compounds have high melting and boiling points
strong electrostatic attraction between anions and cations
57
whats is the electrostatic attraction effected by
charge and distance between particles
58
what happens as you move down group for ionic bonding
elements get larger
59
ioselectronic
equal number of electrons
60
property of isoelectronic ions
larger nuclear charge smaller ion
61
how do predict melting points
look at cations and anions and compare charge and then size
62
why are ionic compounds brittle
when hit opposite charges together and lots of repulsion breaks layers apart
63
how does a covalent bond form between two atoms
shared pare of electrons between atoms
64
how much covalent bonds to atoms form
until they have full outer shell
65
problems with circle and dot diagrams
cant see which electron came from where messy
66
bonding pairs
pairs of electrons bonding in covalent bonds
67
lone pairs
outside unbonded electrons
68
double bonds
2 pairs of electrons shared between 2 atoms
69
triple bonds
3 electrons shared between 2 atoms
70
how to predict which covalent bond will be longer
compare size of atoms forming bonds bigger atoms is longer bonds
71
how to predict melting and boiling point of bonds
compare bond lengths shorter lengths stronger
72
dative covalent bonds
one atom contributes both electrons to the shared pair
73
how to represent dative bond
arrow from which one have it to where its going or two of same dot or cross
74
2 elements to form electrons
dative bond donors and acceptors
75
valence bond theory
if two atoms have orbitals which have unpaired electrons then orbitals will overlap with each other forming a covalent bond
76
how do orbitals overlap
one pair of orbitals overlap head on to form one bond and other bond side on
77
sigma bonds
when orbitals overlap head on lower case sigma
78
pi bonds
orbitals overlap side on
79
why does carbon form 4 bonds instead of 2
ground state = lowest energy orbital in ground state can have 4 unpaired electrons
80
electronegativity
ability to attract electron in covalent bond
81
electronegativity across period
increases as more protons
82
electronegativity down group
decreases down group shell increases so nucleur radius bigger so less atraction
83
electronegativity
ability to attract electron pair in a covalent bond
84
factors that affect electronegativity of bonded pair
-nuclear charge -distance from nucleus -shielding
85
exceptions of trends of electronegativity
-hydrogen similar to C as no electron shielding -dont apply to group 8 -rare bonds
86
dipole
one part of molecule slightly negative and the other slightly positive
87
polar
slightly negative and positive parts of molecule
88
comparing polarity
look at pauling scale values and compare difference
89
ionic compounds formed when electronegativity high
when difference between to atoms so high the electrons almost belong to that atom -atom able to conduct electricity and metal and brittle -electrons sent to other atom
90
metals formed in electronegavitity
when difference lower than 1.5 or greater
91
what holds polyatomic ionic bonds together
covalent bonds
92
ionic compounds with polyatomic ions
covalent and ionic bonds
93
accounting for charge when drawing polyatomic ions
if compound has negative charge add electrons to most electronegative compound -if molecule positive add electron to least electronegative molecule
94

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