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atomic structure Flashcards

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1
Q

describe Dalton’s atomic theory

A

-atoms=tiny particles made of elements
-atoms cannot be divided
-all atoms in an element are the same
-atoms of one element are different to those of other elements

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2
Q

what did Thompson discover about electrons

A

-negative charge
-can be deflected by electromagnetic fields
-very small mass

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3
Q

explain current model of the atom

A

-protons & neutrons found in nucleus
-electrons orbit nucleus in shell
-nucleus=tiny compared to total vol of atom
-most of atom’s mass is in the nucleus
-most of atom=empty space beween nucleus and electrons

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4
Q

charge of proton and electron

A

electron= -1
proton= +1

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5
Q

which particle has the same mass as a proton

A

neutron

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6
Q

which particles make up most of atom’s mass

A

protons and neutrons

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7
Q

what does atomic number tell us about element

A

number of protons in an atom

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8
Q

how is mass number calculated

A

mass number= protons + neutrons

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9
Q

isotope definition

A

atoms of the same element with different number of neutrons (and mass number) but same number of protons (atomic number)

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10
Q

why do different isotopes of the same element react in the same way.

A

-neutrons have no impact on chemical activity
-reactions involve electrons, isotopes have the same electronic configeration

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11
Q

why do different isotopes of the same element react in the same way.

A

-neutrons have no impact on chemical activity
-reactions involve electrons, isotopes have the same electronic configuration

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12
Q

define relative atomic mass

A

weighted mass of an atom of an element compared with one twelfth of the mass of an atom of carbon-12

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13
Q

define relative isotopic mass

A

mass of an atom of an isotope compared with one twelfth of the mass of an atom of carbon-12.

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14
Q

relative isotopic mass is the same as which number

A

mass number

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15
Q

what 2 assumptions are made when calculating mass number

A

-contribution of electron is neglected
-mass of both proton and neutron is taken as 1.0 u

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16
Q

3 uses of mass spectrometry

A

-identify unknown compounds
-find relative abundance of each isotope of an element
-determine structural information

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17
Q

what does the principal quantum number indicate

A

shell occupied by the electrons

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18
Q

shell definition

A

group of orbitals with the same principal quantum number

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19
Q

how many electrons can the 1st shell hold

A

2

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20
Q

how many electrons can the 2nd shell hold

A

8

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21
Q

how many electrons can the 3rd shell hold

A

18

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22
Q

how many electrons can the 4th shell hold

A

32

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23
Q

what is equation for number of electrons in a shell

A

2n²

24
Q

what is an orbital

A

a region around the nucleus that can hold up to 2 electrons with opposite spins

25
Q

how many electrons can an orbital hold

A

2

26
Q

what are the 4 types of orbital

A

S P D F

27
Q

shape of S and P orbital

A

S=spherical
P=dumbell/8 and rotates on axis

28
Q

how many orbitals are found in a S subshell

A

1

29
Q

how many electrons can be held in a S subshell

A

2

30
Q

how many orbitals does P subshell have

A

3

31
Q

how many electrons can be held in a P subshell

A

6

32
Q

how many orbitals are present in a D subshell

A

5

33
Q

how many electrons can be held in a D subshell

A

10

34
Q

how many orbitals are present in a F subshell

A

7

35
Q

how many electrons can fill F subshell

A

14

36
Q

when using ‘electrons in box’ representation, what shape is used to represent electrons

A

arrow

37
Q

what letter is used to represent shell number

A

n

38
Q

from which shell onwards is S orbital present

A

n=1

39
Q

from which shell onwards is P orbital present

A

n=2

40
Q

from which shell onwards is D orbital present

A

n=3

41
Q

from which shell onwards is F orbital present

A

n=4

42
Q

what are the 5 rules by which electrons are arranged in the shell

A

-electrons are added one at a time
-lowest available energy level filled first
-each energy level must be filled before next can be
-each orbital is filled singly before pairing
-4s is filled before 3d

43
Q

why does 4s orbital fill before 3d orbital

A

4s orbital has a lower energy than 3d before it is filled

44
Q

electronic configuration of krypton (Ar=36)

A

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

45
Q

how can the electron configuration be written in short

A

the noble gas before the element is used to abbreviate
E.g. Li→1s2 2s1 ; Li→(He) 2s1

46
Q

how is the group number related to the number of electrons

A

group number= number of electrons in the outer shell

47
Q

how are elements of the periodic table arranged

A

in order of increasing atomic numbers

48
Q

what is meant by periodicity

A

the repeating trends in chemical and physical properties

49
Q

define first ionisation energy

A

energy required to remove one electron from each atom in one mole of gaseous atoms which become one mole of gaseous ions

50
Q

write an equation for the first ionisation energy of magnesium

A

Mg(g)→Mg+(g) + e-

51
Q

factors affecting ionisation energy

A

-atomic radius
-nuclear charger
-electron shielding

52
Q

why does first ionisation energy decrease between group 2 to 3

A

-in group 3 the outermost electrons are in P orbitals
-in group 2 the outermost electrons are in S orbitals so the electrons are more easily removed

53
Q

why does first ionisation energy decrease between group 5 to 6

A

-group 5 electrons in P orbital which are single electrons
-group 6 the outermost electrons are spin paired

53
Q

why does first ionisation energy decrease between group 5 to 6

A

-group 5 electrons in P orbital which are single electrons
-group 6 the outermost electrons are spin paired

54
Q

does first ionisation increase or decrease between end of one period and start of the next

A

decrease
increase in atomic radius=increase in electron shielding

55
Q

does first ionisation increase or decrease down a group

A

-decrease
-shielding increases → weaker attraction
-atomic radius increases → distance between the outer electrons and nucleus increases → weaker attraction
-increase in n. protons is outweighed by increase in distance and shielding

56
Q

describe the structure, forces and bonding in every element across period 2

A
  • Li & Be→ giant metallic; strong attraction between positive ions and delocalised electrons ; metallic bonding
    -B & C → gaint covalent ; strong forces between atoms ; covalent
    -N2, O2, F2, Ne→ simple molecular ; weak intermolecular forces between molecules ; covalent bonding within molecules and intermolecular forces between molecules

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