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Chemistry > Module 2,electrons, Bonding And Structure > Flashcards

Module 2,electrons, Bonding And Structure Flashcards

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1
Q

Compound

A

A substance that is composed of 2 or more different elements

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2
Q

Molecule

A

A group of atoms bonded by the covalent bonds, representing the smallest fundamental unit of a chemical compound that can take part in a chemical reaction. Can be element or a compound.

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3
Q

Ion

A

An atom or molecule with a net charge due to loss/gain of electrons

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4
Q

Molecular ion

A

An ion made up of a group of atoms with an overall charge(usually by gaining/losing a monatomic ion like H+)

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5
Q

Monoatomic ion

A

Formed by gain/loss of electrons in a single atom e.g cl-

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6
Q

Iodide

A

I-

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7
Q

Ammonium

A

NH4+

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8
Q

Hydroxide

A

OH-

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9
Q

Fluoride

A

F-

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10
Q

Carbonate

A

CO3 2-

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11
Q

Nitrate

A

NO3 -

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12
Q

Sulphate

A

SO4 2-

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13
Q

Lithium ion

A

Li+

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14
Q

Aluminium ion

A

Al3+

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15
Q

Chloride

A

Cl-

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16
Q

Hydrogen carbonate

A

HCO3 -

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17
Q

Calcium ion

A

Ca2+

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18
Q

Copper (ii) ion

A

Cu2+

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19
Q

Hydrogen ion

A

H+

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20
Q

Magnesium ion

A

Mg2+

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21
Q

Silver(i) ion

A

Ag+

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22
Q

Phosphate

A

PO4 3-

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23
Q

How many electrons in s, p, d, f sub shell

A

S-2 p-6 d-10 f-14

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24
Q

How many electrons in each orbital

A

2

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25
Q

Notation for electron shells

A

3s, 3 represents its in the 3rd shell, s is type of sub shell

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26
Q

Describe electron shells

A

Electrons have fixed energies. Move in shells/energy levels. Each shell has a number, n (principle quantum number), further from nucleus, higher its energy.

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27
Q

Orbital

A

An area in an atom which can hold 2 electrons. Orbitals in different sub shells have different shapes

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28
Q

What is spin

A

A property electrons have, to do with momentum. 2 electrons in same orbital will have opposite spins, one is up the down. Electrons occupy orbitals in same energy level singly first, before they pair up.

29
Q

Why does the 4s shell fill up before 3d

A

4s has lower energy than 3d

30
Q

What is the aufbau principle

A

Electrons fill up lowest energy sub shells first, closest to nucleus.

31
Q

EQ- Complete the electron configuration of a silicon atom

A

1s2 2s2 2p6 3s2 3p2

32
Q

EQ- how many full orbitals are in an atom of sulfur

A

Draw out the arrows in boxes, then count how many have 2 electrons in, the answer is 7

33
Q

Electron arrangement of bromine

A

Look at atomic number, 35, 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5, short structure, look at previous group 8, [Ar]4s2 3d10, then look at structure of that group 8 and add anything to next to [Ar] that is not in Ar for example

34
Q

What determines what element is where in the periodic table

A

N of electrons in outer shell, group principle quantum number, period outer most electron orbital, block

35
Q

What is the shape if p and s sub shells

A

S are spherical, cross with diagonal line and circle in middle. P have dumbbell shapes

36
Q

What is Heisenberg uncertainty principle

A

Cannot know speed of electrons and exact position, orbitals are where most likely found but not definitely

37
Q

Metallic bonding

A

Strong electrostatic attraction between positive metal ions and delocalised electrons

38
Q

Ionic bonding

A

Strong electrostatic attraction between oppositely charged ions

39
Q

Covalent bonding

A

Strong electrostatic attraction between a shared pair of electrons and positive nuclei

40
Q

What are the properties of a metallic structure

A

Malleable, ductile, conducts electricity, high mp/bp

41
Q

What are the properties of a ionic structure

A

High mp/bp-due to strong electrostatic forces, conducts electricity when l and aq not s, soluble

42
Q

What are the properties of simple molecule structure

A

Low mp/bp, can’t conduct electricity

43
Q

What are the properties of a giant covalent lattice

A

Brittle/soft, high mp/bp, can’t conduct electricity

44
Q

What are the properties of diamond

A

Hard/strong, high mp/bp, doesn’t conduct electricity as it uses all 4 electrons

45
Q

Ocet rule

A

Atoms will share or transfer electrons until they have a full outer shell of 8 electrons

46
Q

What are the exceptions to the ocet rule

A

D block elements, p block elements in 3rd row and below, some p blocks can have less than 8

47
Q

Dative covalent bonding

A

Both electrons come from one atom

48
Q

Describe electron density and lone pairs

A

Electrons repel each other. Electrons exist in area of electron density, 2 types bonding and lone(a pair of electrons not involved in bonding). Single bond counts as 1 electron density and double/triple bond counts as 1 area of electron density. Lone pairs count as area of electron density but repel more than a bonding pair

49
Q

Which pair repel the most to least

A

Most- lone and bonding, lone and lone, bonding and bonding-least

50
Q

How to see how many bonding and lone pairs

A

Draw Dot and cross diagram

51
Q

Describe intermolecular forces

A

Forces between molecules, weaker than covalent/ionic/metallic. Determine mp/bp, van der walls(old term for intermolecular forces, don’t use)

52
Q

Describe London forces

A

Induced dipole-dipole forces, electron move constantly and randomly within an atom, uneven distribution of electrons causes instantaneous dipole, dipole induces a dipole in neighbouring molecule. Simplified-electrostatic force of an attraction between s- on one atom and s+ on other atom of another molecule. More electrons in a molecule, greater possible difference in distribution of electrons, stronger induced dipoles, stronger London forces.

53
Q

Why does bp of group 7 increase down group

A

More electrons in a molecule, greater possible difference in distribution of electrons, stronger induced dipoles, stronger London forces.

54
Q

Describe permanent dipole-dipole interactions

A

Only between polar molecules, stronger than London forces.

55
Q

Describe hydrogen bond relevant to intermolecular forces

A

Only when hydrogen is covalently bonded to F,N,O (very electronegative) resulting in a very polar bond. Between H and lone pair in another molecule. Much stronger than London forces or permanent dipole interactions. Molecules able to form hydrogen bonds, can dissolve in water, higher bp/mp.

56
Q

Why is ice less dense than water

A

Due to hydrogen bonds formed, molecules form an open lattice structure

57
Q

EQ- why does ice have a higher melting point than solid ammonia

A

Ice has stronger hydrogen bonds, O has 2 lone pairs, N has 1(look at dot and cross diagram)

58
Q

EQ- Why do at RTP alcohol are liquids whereas alkanes are gases

A

Alkanes have intermolecular forces, alcohols have hydrogen bonds and intermolecular forces, hydrogen bonds stronger

59
Q

EQ- why does NF3 have a permanent dipole

A

Is pyramidal, dipoles do not cancel out, F more electronegative than N

60
Q

EQ- why does SF6 have a octahedral shape

A

because it has 6 bonding and 1 lone, electrons repel as far as they can

61
Q

What to include in a extended answer about shape of molecules

A

Electrons group repel as far as possible, number of electron density, number of bonding/lone pairs, lone pairs repel more than bonding, this causes shape… with bond angle of…

62
Q

Electronegativity

A

A measure of the ability of an atom in a molecule to attract electrons in a chemical bond

63
Q

What 2 factors affect electronegativity

A

Radius-smaller than radius of atom, greater attraction to the bonded pair of electrons(less distance between electrons and positive nuclei, stronger attraction). Nuclear charge-greater nuclear charge, stronger attraction.

64
Q

Describe polar bond

A

2 atoms bonded have different electronegativity’s, has dipole negative and positive side. Has polar bonds and no symmetry or symmetry but doesn’t cancel out charges.

65
Q

Describe non-polar bond

A

2 atoms bonded have same/similar electronegativtys. No polar bonds,or there is symmetry.

66
Q

What is the most electronegative

A

Fluorine

67
Q

What is the order of sub shells

A

1s,2s,2p,3s,3p,4s,3d,4p,4d

68
Q

What does electronegativity mean

A

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

Chemistry (6 decks)

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