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chemistry > 1.Foundations in chemistry, Periodic table and enegy > Flashcards

1.Foundations in chemistry, Periodic table and enegy Flashcards

topic 1 and 2

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1
Q

Ionic bond definition

A

the electrostatic attraction between opp. charged ions

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2
Q

covalent bond definition

A

the strong electrostatic attraction between a SHARED PAIR of electrons and the NUCLEI of the bonded atoms

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3
Q

linear- bond angle /electron pairs/ example

A

180/ 2 bonding pairs of electrons / Co2

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4
Q

trigonal planar- bond angle /electron pairs/ example

A

120/ 3 bonding pairs of electrons/ Bcl3

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5
Q

tetrahedral- bond angle /electron pairs/ example

A

109.5 / 4 bonding pairs of electrons /CH4

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6
Q

pyramidal- bond angle /electron pairs/ example

A

107/ 3 bonding + 1 lone pair/ NH3

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7
Q

non-linear- bond angle /electron pairs/ example

A

104.5 / 2 bonding = 2 lone / H2O

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8
Q

octahedral- bond angle /electron pairs/ example

A

90 / 6 bonding pairs of electrons / SF6

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9
Q

electronegativity definition

A

the ability of an atom to attract the bonding electrons in a covalent bond

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10
Q

first ionisation energy definition

A

the energy required to remove one electron from every atom in one mole of GASEOUS ATOM

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11
Q

metallic bond definition

A

a strong electrostatic attraction between CATIONS and delocalised electrons

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12
Q

4 types of structure and boning

A

giant ionic
giant covalent
giant metallic
simple molecular/ covalent

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13
Q

describe the structure and bonding of a giant ionic lattice

A

-giant structure
-ionic bonding
-strong electrostatic forces of attraction
-in a GIANT IONIC LATTICE where forces act in ALL directions

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14
Q

3 properties of ionic compounds

A

-high melting/ boiling point
-conducts electricity when DISOLVED in water/ moltten
-poor conductors when SOLID

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15
Q

why do giant ionic compounds have a high melting point?

A
  • strong elaectrostatic forces oif attraction IN ALL DIRECTIONS
    -between opp. charged ions
    -STRONG ionic bonds need lots of energy to break
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16
Q

why do giant ionic compounds have good electrical conductivity when molten/ dissolved in water?

A

-when molten/ disolved in water
-ions can m,ove freeley
-ions are mobile charge carriers

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17
Q

why do giant ionic compounds have bad electrical conductivity when solid?

A

-when solid
-ions are fixed
-and cant move
- are not mobile charge carriers

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18
Q

another word for dative covalent bonding?

A

coordinate bonding

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19
Q

why are electrons spread out in pairs?

A

-electron pairs repel ( as both are -ve)
-forced maximum distance apart
-lone pairs repel more than bonded pairs of electrons

(pairs of electrons repel other pairs of electrons)

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20
Q

pattern in electronegativity on the periodic table? and explination?

A

-increases as you move across a period (L-R)

more protons in the nucleus from (L-R)
increase in nuclear attraction between protons in the nucleus and electrons in the outer shells

-increases as you move up a grp

as you move up a grp atom size decreases
it is easier to get an electron closer to the nucleus

(increases in a diagonal across periodic table from FR to F)

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21
Q

what does a delta (δ) symbol show in a permenant dipole?

A

-a slight charge

δ+ is least electronegative
δ- is most electronegative

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22
Q

what is the most electrongative element?

A

flourine

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23
Q

what is a polar molecule?

A
  • have bonds with a permanent dipole
    -the molecular shape must not cause dipoles to cancel each-other out
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24
Q

what are the anomalous properties of water? and explain?

A
  • Ice is less dense then water
    coz ice has an OPEN lattice structure
    Hydrogen bonds in the lattice hold the water molecuels
    apart

-water has a relativley high boiling and melting point
coz it has hyfrogen bonds in addition to london forces
so more energy is needed to overcome the additional hydrogen bonds

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25
Q
A
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26
Q

what 3 elements can a hydrogen bond be formed with?

A

F
O
N

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27
Q

what delta charge is hydrogen

A

δ+

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28
Q

what delta charge does F/O/N have?

A

δ-

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29
Q

why metals have high melting point

A

-giant mettalic lattice
-strong metalic bonding
-lots of energy to break

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30
Q

why metals are good cobductors of electricity

A

have mobile delocalised electrons
which are mobile charge carriers

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31
Q

why metals dont disolve in water?

A

metalic bonds are too strong to be broken down by water

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32
Q

graphite structure

A

giant covalent lattice
layers of hexogons(trigonal planer angle 120)
each carbon atom has 3 strong covalent bonds
1 delocalised electron

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33
Q

graphite properties

A

good conducter of electricity
high melting point
desont disolve in water

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34
Q

why is graphite a good conductor of electricity?

A

each carbon atom has a delocalised electron
which can move through the structure
acting as mobile charge carriers

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35
Q

why graphit has a high melting point

A

giant covalent lattice
3 strong covalent bonds per atom
lots of energy to break

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36
Q

why dosent it disolve in wtaer?

A

not polar
has strong covalent bonds whicha re too strong to be broken down by water

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37
Q

structure of diamond and silicone

A
  • giant covalent lattices
    -each c atom has 4 strong covalent bonds
    (no layers)
    (no delocalised electrons)
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38
Q

diamond and silicone properties?

A

high melting point

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39
Q

why do diamond and silicone have a high melting point?

A

giant covalent lattice
4 strong covalent bonds per c atom
reuires lots of energy to break

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40
Q

why are silicone and graphite poor conductors of electricity

A

no delocalised electrons
which act as mobile charge carriers

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41
Q

nitric acid formula

A

HNO3

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42
Q

difference in melting point in the same period exsplained?

e.g P4 and cl2

A

P4 has a higher melting poit because it has stronger london forces (coz more e-)

more energy needed to break the intermolecular forces

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43
Q

why is a mollecule is polar

A

molecule isnt symetrical
dipoles dont cancel out

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44
Q

giant ionic lattice defenition

A

repeating pattern
of oppositley charged ions

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45
Q

orbital defenition

A

a region around the nucleus that can hold 2 el;ectrons (with opposite spin)

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46
Q

order of filling

A

1s/ 2s/2p/3s/3p/ (4s) /3d/4p

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47
Q

relative atomic mass (RAM) definition

A

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

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48
Q

relative isotopic mass (RIM)

A

the mass of an isotope relative to 1/12th of the mass of an atom of carbon-12

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49
Q

RAM equation

A

(mass x abundance) + (mass x abundance)/100

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50
Q

the 3 general reactions of acids

A

1) Acid + Alkali/base -> salt + water
2) Acid + Carbonate -> salt + carbon dioxide + water
3) acid + metal -> salt+ hydrogen

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51
Q

nutralisation reaction

A

H+ +OH- -> H2O

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52
Q

layers in an electron shell from biggest to smallest

A

shell-> subshell-> orbital

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53
Q

how many electrons in an orbital?

A

2

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54
Q

s subshell orbitals and electtrons?

A

s 1 orbital (2e-)

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55
Q

p subssell orbitals and electrons

A

3 orbitals n(6e-)

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56
Q

d subshell orbitals and electrons?

A

5 orbitals (10e-)

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57
Q

f orbitals

A

7 orbitals (14e-)

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58
Q

p orbital shape

A

dumbell

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58
Q

s orbital shape?

A

sphere

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59
Q

isotope definition

A

atoms with the same number of protons but a different number of nutrons and different masses.

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60
Q

what is the relative mass of an electron

A

1/1836

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61
Q

what are charges and masses of subatomic particles relative to

A

the proton

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62
Q

short hand electron arangement

A
63
Q

isotopes of hydrogen

A
64
Q

acid deffenition

A

H+ proton donor that releases H+ ions in aqueous solution

65
Q

what 4 acids do u need to know?

A

1) hydrochloric acid -> HCL
2)Nitric Acid -> HNO3
3)H2SO4 -> Sulphuric Acid
4)ethanoic acid -> CH3COOH

66
Q

alkali definition

A

soluble basses that release OH- ions in aqueous soloution

67
Q

4 alkalis to know

A

1) sodium hydroxide ->NAOH
2)potassium hydroxide
->KOH
3)amonia->NH3
4)amonia in water (amonium hydroxide)->NH4OH

68
Q

what is a polyatomic ion

A

an ion contyaining 2 or more eleements

69
Q

what is Avogadro’s constant

A

6.02×10²³ mol⁻¹.

70
Q

STRONG AND WEAK ACIDS DISSOCIATIONS

A
71
Q

moles formula with Avogadro’s constant ?

A

N number of atoms
——————— = ————————————–
n x Av moles x Avogadro’s constant

72
Q

moles calculation and steps for solids?

A

steps
1)calculate moles
2)the ratio
3)mass

                 MASS
      --------------------------
       MOLES  I    Mr
                      I
73
Q

what are gasses measured in?

A

volume

74
Q

what are the calculations for moles in gasses? AND THE STEPS

A

volume (cm3) volume (dm3)
——————————————- or ——————————–
moles I 24, 000(cm3) moles I 24(dm3)
I I

steps
1)calculate moles
2)the ratio
3)voleme

75
Q

what is mole of any gas?

A

24,000 cm3 or 24 dm 3

76
Q

moles calculations for solutions (aq)?

A

conc. I volume

77
Q

what is the unit for concentration

A

mol/dm3

78
Q

titrations

A
79
Q

what is the empirical formulae?

A

the simplest whole number ratio of atoms in a compound

80
Q

steps for emirical formulae

A

steps
1)mass
2) mass/Ar
3)divide all ratios by the lowest ratio

81
Q

water of crystalisation

A
82
Q

Ideal gas equation and units

A

pV = nRT

P= pressure in pascal (Kpa x 1000= Pa)

volume =m3 (dm3–>m3 is ÷ 1000) (cm3 -> m3 is ÷ 1,000,000)

n= moles

R= gass constant in exam aid (8.3-4 J mol-1 K-1)

T= temperature in kalvin (K)

83
Q

formula for Kalvin

A

C+ 273=K

c= Celsius

84
Q

formula for celcius

A

k-273= Celsius

k= Kalvin

85
Q

what is ideal gass?

A

1 mole of any gas - occupies 24dm3 at room temp and pressure

86
Q

what temp. is room temp.?

A

20 degree c or 293k

87
Q

what is room pressure

A

101 kpa

88
Q

how to reduce percentage uncertanty in mass

A
  • use balance that weighs to more decimal places
  • use a larger mass
89
Q

use of aqueous barium chloride in quantitative analysis test

A

-test for sulphate
-white precipitate forms

90
Q

-test for sulphate ions and results

A

use of aqueous barium chloride in quantitative analysis test

-white precipitate forms

91
Q

molecular formula deffenition

A

the actual number of atoms of each element in a compound

92
Q

Q
what is the oxidation state of an element?

A

0

93
Q

why increased temp. means increased rate of reaction

A

-increased temp. means activation energy is reached

-so increased kinnetic energy of reactant particles
-increased frequency of succesfull collisions between reactant particles

94
Q

draw catalyst and boltzmann
and explanation

A

-catalyst decreases the EA
-by providing an alternative reaction pathway

-so a greater proportion of molecules have energy exceedin EA

-causes an increased rate of reaction

95
Q

draw temperature and boltzmann
and explanationv

A

-increased temp
-so greater proportion of molecules with energy exceeding the EA
- increased kinnetic energy so more FREQUENT collisions

-therfore increased rate of reaction

96
Q

rules for drawing energy diagram

A
  • draw all the arrows facing up
    then go back and draw arrows facing down
97
Q

oxidation in terms of electrons

A

loss of electrons

98
Q

reduction in terms of electrons

A

gain of electrons

99
Q

oxidation in terms of oxidation state

A

increase in oxidation state
e.g (0 to +2)

100
Q

reduction in terms of oxidation state

A

reduction in oxidation state (0 to -2)

101
Q

a general equation for a typical redox reaction

A

metal+ acid -> salt+ hydrogen

102
Q

what is calcium hydroxide

A

lime water

103
Q

hydroxide general equation

A

oxide(s) + water(l) → hydroxide(aq)

104
Q

what causes an increase in ph (not spec friendly)

A

more hydroxide ions, OH–, dissociate into a solution cause the pH to increase

105
Q

trend in solubility of group 2

A
  • increase solubility down the grp
    -higher ph/ becomes more alkaline
    -coz more OH- released
    -so there is a higher conc. of oh- ions present.
106
Q

barium oxide and water equation

(BARIUM CAN BE SWAPPED WITH LITERALLY ANY GRP 2 METAL)

A

BaO(s) + H₂O(l) →Ba(OH)₂ (aq)

107
Q

another name for group 2

A

alkaline earth metals.

108
Q

what are the uses for group 2 Compounds

A
  • Indigestion treatments
  • agriculture
109
Q

what grp 2 compounds are used in agriculture and why

A

calcium hydroxide Ca(OH)₂

to increase the Ph of the soil and make it more alkali

equation Ca(OH)₂ + H₂SO₄ → CaSO₄ + 2H₂O

110
Q

what grp 2 compounds are used in indigestion treatment

A

CaCO3 EQUATION: CaCO₃ + 2HCl → CaCl₂ + H2O + CO₂

and

Mg(OH)₂ EQUATION: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O +

111
Q

why are group 2 metals s block elements

(why is nay element in any block)

A

the highest energy electron is in an s sub shell

112
Q

what is the trend in reactivity of group 2 and why

A

-the reactivity of grp 2 increases DOWN a period
-and become more powerful reducing agents
-because the outer shell electron is lost more easily
COZ
-sum of 1st and 2nd ionisation energy decrease down the grp
-nuclear attraction decreases
- coz larger atomic radius means greater electron sheilding

113
Q

what are the 3 redox reactiions of grp 2 that I need to know?

A

Metal+water -> metal hydroxide + hydrogen

Metal+oxygen -> Mrtal oxide

Metal+acid ->salt + hydrogen

114
Q

general equation for grp 2 and oxygen

A

2M(s) + O₂(g) → 2MO(s)

115
Q

general equation for grp 2 and water

A

M(s) + 2H₂O(l) → M(OH)₂ (aq) + H₂

116
Q

general word equation for grp 2 and acids

A

metal + acid → salt(aq) + hydrogen(g)

117
Q

what shell do all halogens end in

A

(s2 or p5)

118
Q

what is the grp 7 trend in boiling point

A
  • boiling point increases down the group
    -more electrons
    -more London forces between molecules
  • more energy needed to overcome
119
Q

what is the grp 7 trend in reactivity

A
  • decrease down a grp
  • weaker oxidising power
    -increased number of electrons
    -more shielding
    -larger atomic radius
    -weaker nuclear attraction
    -harder to gain electrons
120
Q

Observations with an organic solvent

A
121
Q

Observations in aqueous solution

A
122
Q

what are the 3 displacement reactions I need to know

A

Cl and Br
Cl and I
Br and I

123
Q

what are the 3 displacement reactions Ionic I need to know

A

1) Cl₂ (aq) + 2Br-(aq) → 2Cl- (aq) + Br₂ (aq)
2) Cl₂ (aq) + 2I-(aq) → 2Cl- (aq) + I₂ (aq)
3) Br₂ (aq) + 2I-(aq) → 2Br- (aq) + I₂ (aq)

124
Q

what is disproportionation

A

when the same element is oxidised and reduced

125
Q

what example of a disproportionation do i need to learn

A

chlorine with COLD aqueous sodium hydroxide

Cl₂(aq)+ 2NaOH(aq) → NaCl(aq)+ NaClO(aq) + H₂O (l)

126
Q

NaClO(aq) household name

A

bleach

127
Q

bleach chemical formula

A

NaClO(aq)

128
Q

halides with silver nitrate general ionic equation

A

Ag+ (aq) + X-(aq) → AgX(s)

129
Q

what precipitates form with silver nitrate chlorine bromine and iodine

A

AgCl = white

AgBr= cream

AgI+ yellow

130
Q

further testing with amonia results

A

AgCl = redisolves in dilute amonia

AgBr= redissolves in concentrated amonia soloution

AgI+ does not redisolve

131
Q

Homologous series

A

a series of organic compounds having the same funtional group but with each sucessive members differing by CH2

132
Q

what are the blocks in the periodic table

A

s,d,p,f

132
Q

divide a periodic table into its blocks

A
133
Q

fraw 2 molecules of water together

A
133
Q

trend in electronegativity

A

-increases across a grp
(as atomic radius decreases more protons in nucleus as u move from l-R)

decreases as you move down a grp

coz atom size increases so it is harder to get an electron closer to the nucleus

133
Q

what 3 elements form a hydrogen bond

A

f
o
n

134
Q

anomolous prperties of water

A

-ice is less dense than water
-open lattice structure
-the H bonds in the lattice hold the water molecules apart

-water has a relativley high boiling point
-coz it has H bonds in addition to london forces
-which require lots of energy to overcome

134
Q

1st and 2nd ionisation energy equation

A

E(g) –> E+ + e-

E+(g)–> E2+ +e-

134
Q

why does ionisation energy increase evrytime you remove an electron

A

because increased +ve charge

so greater nuclear attraction by remaining electrons

135
Q

Ionisation energy down a grp

A
136
Q

Ionisation enegy across a grp`1

A
137
Q
A
138
Q
A
138
Q
A
138
Q
A
138
Q
A
139
Q
A
139
Q
A
140
Q
A
140
Q
A
141
Q
A
142
Q
A
143
Q
A
143
Q
A
144
Q
A
144
Q
A

chemistry (5 decks)

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