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Chemistry > Structure and Bonding - Week 1-6 > Flashcards

Structure and Bonding - Week 1-6 Flashcards

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1
Q

What is the mass of an atom?

A

the mass of the nucleus

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2
Q

What makes up the nucleus?

A

Protons and neutrons

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3
Q

What is the relationship between Daltons and g mol^-1?

A

1 Da = 1 g mol^-1

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4
Q

What is the Avogadro constant?

A

Na = 6.02214x10^23 mol^-1

Defines how many (atoms/molecules/things) there are in a mole

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5
Q

What is the unit for atomic mass?

A

Daltons (Da)

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6
Q

What is charge measured in?

A

Coulombs (C) in SI units

Protons +1
Electrons -1

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7
Q

How is nuclear charge determined?

A

number of protons in the nucleus

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8
Q

What is the mass number?

A

Number of protons + electrons

(top number)

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9
Q

What 4 factors affect electrons within atoms?

A

Energy level
Orbital shape
Orbital direction
Spin

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10
Q

How are electrons energy levels grouped and shown?

A

Energy is quantised according to number of nodes (wiggles present in the electrons wave)

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11
Q

What does the number of nodes represent?

A

energy level

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12
Q

What does it mean when an electron has more nodes?

A

more energy

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13
Q

What are the 3 types of orbital shapes?

A

s orbital
p orbital
d orbital

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14
Q

How many nodes are present in a s orbital?

A

no visible nodes present

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15
Q

How many nodes are present in a p orbital?

A

one node visible

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16
Q

How many nodes are present in a d orbital?

A

2 nodes visible

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17
Q

How many orbitals are allowed in the s orbital and what is their direction?

A

Spherical (not directional)

Only one s orbital per energy level

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18
Q

How many orbitals are allowed in the p orbital and what is their direction?

A

Three orbitals allowed per energy level

Point along x, y, z axes

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19
Q

How many orbitals are allowed in the d orbital?

A

5 orbitals allowed per energy level

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20
Q

Which way do electrons spin?

A

Internal spin
Only one value of spin is allowed
Electrons can spin in two different directions
-> spin up
-> spin down

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21
Q

What are the 3 electron filling rules?

A

Aufbau rule - electrons occupy the lowest energy orbital
available
Pauli principle - maximum of 2 electrons per orbital - one
spin up the other down
Hund’s rule - for orbitals with the same energy orbitals
are occupied singly if possible - spreads
electrons out

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22
Q

Why do unstable nuclei undergo nuclear decay?

A

to become more stable

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23
Q

What is alpha decay?

A

the nucleus will eject a particle containing 2 protons and 2 neutrons

4 He
2 —> alpha particle

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24
Q

What is beta decay?

A

a neutron in an atom can split into a proton and an electron
proton stays in the nucleus but an electron leaves (emitted electron = beta particle)

0 e
-1 –> B particle

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25
Q

What is nuclear fission?

A

Nuclei can react by absorbing neutrons and splitting into two smaller nuclei.

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26
Q

What is nuclear fusion?

A

Nuclei of small atoms can combine to produce a larger nucleus.

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27
Q

What is ionic bonding?

A
  • one atom loses and electron to form a
    positively charged cation
  • one atom gains an electron to from a
    negatively charged anion
  • cations and anions are held by attraction
    between +ve and -ve charges
28
Q

What is needed for ionic bonding to occur?

A

One element good at loosing electrons
- low electronegativity
- a metal
One element good at gaining electrons
- high electronegativity
- a non-metal

29
Q

What is covalent bonding?

A

two atoms sharing a pair of electrons

the electron pair will be exactly half way between the atoms when the electronegativity of each atom is identical

30
Q

What is needed for a bond to be 100% covalent?

A

both atoms must have identical electronegativity

e.g. H-H, Cl-Cl

31
Q

What is a polar covalent bond?

A

a bond made up of elements with different electronegativities
e.g. Si-O

32
Q

What occurs with charges in a polar covalent bond?

A
  • a dipole is created

Because the electrons are not evenly shared
one becomes slightly negative (delta negative)
one becomes slightly positive (delta positive)

33
Q

What happens to the overlap of wavefunctions for two 1s orbitals approaching each other in phase (same sign)?

A

Overlap is constructive

34
Q

What happens to the overlap of wavefunctions for two 1s orbitals approaching each other out of phase (opposite sign)?

A

Overlap is destructive

35
Q

What does a constructive overlap result in?

A

= bonding

Electron density between atoms increases

36
Q

What does a destructive overlap result in?

A

= Anti-bonding

Electron density between atoms decreases

37
Q

What happens to the extent of orbital overlap as a result of separation?

A

Extent of orbital overlap increases as separation between orbitals decreases.

38
Q

What are the 3 orbital overlap rules?

A
  1. only orbitals with similar energies can
    overlap
  2. for every bonding orbital created a higher
    energy anti-bonding orbital is also created
  3. orbital symmetries must match for overlap
    to occur
39
Q

What are the orbital interactions of sigma orbitals for s orbitals?

A

interactions are head on - resulting in bonding / antibonding orbitals forming from s orbitals

40
Q

What are the orbital interactions of sigma orbitals for p orbitals?

A

p orbitals can overlap head on

41
Q

What are the orbital interactions of p orbitals that from pi-bonds ?

A

Side on overlap

42
Q

How can non bonding electrons occur?

A

Sometimes neither a bond or anti-bond can form due to mismatch in number of orbitals

43
Q

How does electron shielding differ with orbitals?

A
  • Proportion (%) of shielding electrons
    decreases across a row - strength that
    electrons are held increases

s orbitals > p orbitals > d orbitals

44
Q

In the periodic table what do rows and columns show?

A

Rows - filling of outermost electron energy
levels
Columns - group elements with identical
arrangement for their outermost
electrons

45
Q

What are valence electrons?

A

The outermost electrons in an atom
The only electrons involved in bonding atoms together to make molecules

46
Q

What are core electrons?

A

Electrons never involved in bonding - too deep within atom to interact with electrons from other atoms

47
Q

What is an example of the electron configuration of Uranium?

A

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f3 6d1

(can abbreviate electrons using configurations of group 18 elements)

48
Q

Characteristics of metals?

A
  • not good at looking after valence electrons
  • lose valence electrons to form positively
    charged cations
  • conduct heat and electricity well
  • shiny
49
Q

What is the rule for d-block metal’s electron count?

A

d-electron count = group number - metal charge

50
Q

Characteristics of non-metals?

A
  • good at looking after valence electrons -
    difficult to remove from non-metals
  • can absorb extra valence electrons forming
    negatively charges anions
51
Q

What causes the difference in metal and non-metal’s ability to hold on to valence electrons?

A

Shielding = elements on the left hand side of periodic table hold valence electrons less strongly than those on the right hand side

52
Q

characteristics of non-metals near to the metal/non-metal boundary?

A
  • hold valence electrons quite strongly
  • not great at attracting electrons from other
    species
  • not good at making cations or anions
53
Q

Characteristics of atoms within groups?

A

Similar characteristics
Electrons easier to remove as you go down the group - outer electrons get further from nucleus

54
Q

Characteristics of unreactive elements?

A

Elements with a completely filled valence energy level.

  • Difficult to remove electrons
  • Difficult to gain electrons
  • Difficult to share electrons
55
Q

What is the definition of electronegativity?

A

Ability of an atom to attract another electron towards itself when interacting with another atom.

Increases from left to right.
Decreases from top to bottom.

56
Q

What are Lewis structures?

A

Predict covalent bonding between atoms

Represent covalent bonds between atoms as sticks (2 electrons)

57
Q

What are the 6 key steps to drawing Lewis structures?

A
  1. Count all valence electrons
  2. Decide on central atom
  3. Assemble bonding framework with single
    bonds
  4. Place 3 pairs of electrons on each outer
    non-hydrogen atom
  5. assign remaining valence electrons to
    central atom
  6. Minimise formal charges
58
Q

Which direction do wedges represent bonds going?

A

Wedges = bonds coming towards us - out the
paper

59
Q

Which direction do dashes represent bonds going?

A

Dashes = bonds going away from us

60
Q

What are the 6 types of 3D shape arrangement?

A

Linear, Trigonal, Tetrahedral, Square planar, Trigonal bipyramidal, Octahedral

61
Q

What are the 5 steps for predicting shape by Valence Shell Electron Pair Repulsion? (VESPR)

A
  1. Count central atom valence electrons
  2. Add electrons used by other atoms to make
    bonds to central atom
  3. Adjust for charge
  4. Adjust for multiple bonding
  5. Number of pairs predicts shape
62
Q

What is hybridisation of atomic orbitals?

A

Mixing atomic orbitals together.
Orbitals point towards the atoms making the bonds.

63
Q

What molecules are sp hybridised?

A

Linear molecules

  • 2 hybridised orbitals at 180 degrees
  • p orbitals point along axes
  • s orbitals have no direction
64
Q

What molecules are sp2 hypridised?

A

Trigonal planar molecules

  • 3 hybridised orbitals at 120 degrees to each
    other
  • a s and 2 p orbitals combined
  • one p orbital unaffected
65
Q

What molecules are sp3 hybridised?

A

Tetrahedral molecules

  • all available valence orbitals are combined
66
Q

What is the shape and hybridisation of molecules adopted by the valence shell electron pairs (VSEP)?

A

VSEP Shape Hybridisation

2 Linear sp
3 Trigonal sp2
4 Tetrahedral sp3
4 Square planar sp2d
5 Trigonal bipyramidal sp3d
6 Octahedral sp3d2

Chemistry (10 decks)

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