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chemistry > topic 1 > Flashcards

topic 1 Flashcards

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1
Q

what is the structure of an atom

A

nucleus- protons and neutrons
electrons - orbit nucleus

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2
Q

what are the relative charges and masses of the sub atomic particles

A

relative charge
proton - 1
neurton - 0
electron - -1

relative mass
proton - 1
neutrons - 1
electrons - 1/2000

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3
Q

wwhat is the mass number and where is it on the periodic table

A

top number
protons and neutrons

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4
Q

what is the atomic number and where is it on the periodic table

A

bottom number
protons

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5
Q

what are ions

A

have a different number of electrons and protons
negative ions have gained electrons
positive ions have lost electrons

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6
Q

what is an isotope

A

elements with same number of protons but different number of neutrons

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7
Q

what is relative atomic mass

A

the 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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8
Q

what is relative isotopic mass

A

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

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9
Q

what is relative molecular mass

A

the mean mass of a molecule compared to 1/12th of the mass of an atom of carbon-12

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10
Q

how to work out relative atomic mass

A

total abundance

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11
Q

how to calculate isotopic mass

A

place x in isotopic mass you dont know
(abundance A X m/zA)+(abundance B X m/z B)
———————————
total abundance

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12
Q

how to predict mass spectra from % abundance

A
  1. write % as decimals
  2. create a table with isotope combinations
    multiply decimal form abundances to get relative abundance from each molecule
  3. any molecules which are the same add the aundances
  4. divide the relative abundances worked out by the smallest value to give whole number ratio to put on spectra
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13
Q

what is M+1 peak

A

molecular peak ion
same as Mr

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14
Q

what are the subshells with number of orbitals and number of electrons

A

s- 1 orbital 2 electrons
p- 3 orbitals 6 electrons
d- 5 orbitals 10 electrons
f- 7 orbitals 14 electrons

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15
Q

what subshells do shells 1 2 and 3 have and how many electrons can they hold

A

shell 1 - 1s- 2 electrons
shell 2 - 2s 2p- 8 electrons
shell 3 - 3s 3p 3d - 18 electrons

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16
Q

what is the shape of a s orbital

A

spherical

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17
Q

what is the shape of a p orbital

A

dumbells

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18
Q

what is spin pairing

A

when 2 electrons occupy 1 orbital they spin in opposite directions

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19
Q

how do electrons fill orbitalsand why

A

singly first the pair up due to electron repulsion

20
Q

when are subshells most stable

A

when full or half full

21
Q

where are the s d p and f blocks on the periodic table

A

s - group 1 and 2
d- transition metals
p - group 5 6 7 and 0
f - under d block

22
Q

what is the electromagnetic spectrum

A

radiowaves
microwaves
infrared
vissible light
ultra violet
x rays
gamma rays

23
Q

which side of the electromagnectic has the largest energy and frequcency of radiation

A

gamma rays side

24
Q

what is line spectra evidence for and how

A

quantum shells
energy levels have discrete values and electrons can move from one shell to another
a line spectra shows the frequency of light given when a electron moves down

25
Q

on line spectra where has the electron dropped to if lines are at ultraviolet

A

ground state

26
Q

on line spectra where has the electron dropped to if lines are at visible light

A

n = 2

27
Q

one line spectra where has the electron dropped to if lines are at infrared

A

n = 3

28
Q

why is emission spectra evidence for quantum shells

A

they are diffined lines prove electrons exist in shells only and never between them

29
Q

what is ionisation energy

A

min amount of energy required to remove 1 mole of electrons from 1 mole of atoms in gaseous state

30
Q

what is the equation for the 1st ionisation energy Na

A

Na (g) –> Na^+ (g) + e^-

31
Q

what affects the 1st ionisation energy

A

shielding - more shells between positive nucleus and negative electrons means weaker attraction = less energy
nuclear charge - more protons in nucleus the greater the attraction between outer electrons = more energy
atomic size - bigger the atom the further the outer electrons, attractive force reduces = less energy

32
Q

what happens to ionisation energy down group and why

A

decreases
atomic radius increases
sheilding increases

33
Q

what is successive ionisation energy

A

the removal of more than 1 electrons from the same atom

34
Q

draw a graph a the successive ionisation of Mg and explain

A

gerneral increase - more positive ion

large increase after 1st 2 electrons as removing from a shell closer to nucleus

large increase after next 8 as new shell

35
Q

what happens to atomic radii across a period
amd why

A

decreases across
increased nuclear charge- more protons pulls outer shell towards nucleus
extra electrons go into same shell

36
Q

what is the trend in 1st ionisation energy across a period and why

A

increases
increased nuclear charge
similar sheilding

37
Q

draw a graph to show 1st ionisation energy from Na to Ar

A

see flash card

38
Q

what does the fact the 1st ionisation energy of Al lower than Mg provide evidence for

A

subshells
3p1
higher enrgy subshell further from nucleus

39
Q

what does the fact that the 1st ionisation energy of S being similar to P provide evidence for

A

for electron repulsion
3p4

40
Q

draw a graph for the melting points of Na ro Ar

A

see flash card

41
Q

explain the bp of Na Mg and Al

A

metalic bonding
general increase as metal ions have increasing positive charge and delocalised electrons and have smaller ionic radius

42
Q

explain the mp of Si

A

giant covalent
many covalent bonds

43
Q

explain the mp of P

A

weaker simple molecular structure
weak london forces

44
Q

expkain the mp of S

A

larger simple molecular structure
larger london forces

45
Q

explain the mp of Cl

A

simple mocleular structure
smaller london forces

46
Q

explain the mp of Ar

A

exists as individual atoms
smaller london forces

chemistry (15 decks)

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