Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

Chemistry Year 1 > Module 2 (Foundations In Chemistry) > Flashcards

Module 2 (Foundations In Chemistry) Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
1
Q

What is the relative mass and charge of a proton?

A

-Relative Mass = 1.00
-Charge = 1+

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

What is the relative mass and charge of a neutron?

A

-Relative Mass = 1.00
-Charge = 0

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

What is the relative mass and charge of an electron?

A

-Relative Mass = 1/2000
-Charge = 1-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

How to calculate mass number?

A

The total number of protons and neutrons in the nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What is an Isotope?

A

Isotopes are atoms of the same element with different numbers of neutrons as such a different mass number, they have the same atomic number

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

What is relative is relative isotomic mass?

A

Relative isotopic mass is defined as: mass of an atom of the isotope compared with 1/12th of the mass of an atom of carbon-12.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

What is Relative Atomic Mass?

A

The weighted mean mass of all isotopes of the same element within a sample compared with 1/12th of the mass of an atom of carbon-12.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

What is Relative Molecular Mass, Mr?

A

The sum of all the atomic masses present with in a simple covalent molecule.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

What is Relative Formula Mass?

A

The sum of the atomic masses of all the atoms in the formula of the substance (giant structure)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

What are Mass Spectrometers used for?

A

Mass spectrometers can be used to find the abundance of different isotopes present in a sample of an element. The results obtained can be used to calculate the relative atomic mass of the sample.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

What is the Formula for:
-A nitrate ion
-A carbonate ion
-A sulfate ion
-A hydroxide ion
-An ammonium ion
-A zinc ion
-A silver ion

A
  • NO3^-
  • CO3^2-
  • SO4^2-
  • OH^-
  • NH4^+
  • Zn^2+
  • Ag^+
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

How to predict charges on ions?

A

-Group 1 = 1+
-Group 2 = 2+
-Group 3 = 3+
-Group 5 = 3-
-Group 6 = 2-
-Group 7 = 1-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

How to measure amount of substance?

A

The amount of substance is measured using moles (mol). One mole of any substancecontains the same number of particles as there are carbon atoms in 12.00g of Carbon-12.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

How to calculate number of moles?

A

Mol = mass(g) / molar mass (gmol^-1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

What is Avagadro’s constant?

A

The number of atoms in one mole of carbon-12 which is 6.02x10^23, this value is avagadro’s constant it can also be applied to other particles.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

How to calculate number of particles?

A

number of particles = n (mol) x 6.02x10^23

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

How to calculate volume of gas?

A

Volume of gas (dm^3) = n (mol) x 24

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

What id the Ideal Gas Equation?

A

pV = nRT

p = Pressure (Pa)
V = Volume (m^3)
n = Moles (mol)
R = Gas constant (8.314)
T = Temperature (K)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

What are the Two Types of Formuale?

A

-Mollecular (full formulae)
-Empirical (simplest ratio)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

What is the equation to calculate concentration?

A

concentration (moldm^-3) = n (mol) / V (dm^-3)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Formula for Percentage yield

A

%Yield = actual amount (mol) of product / theoretical amount (mol) of product x 100

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

Formula for Atom economy

A

Mr of desired product / Mr of all products x 100

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

What is an acid?

A

A compound that disociates when dissolved in water to release H+ ions. They are commonly known as proton donors

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

What Common acids do you need to know?

A

-Hydrochloric Acid (HCl)
-Sulfuric Acid (H2SO4)
-Nitric Acid (HNO3)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

What Makes an Acid Strong?

A

The ability to fully dissociate into H+ ions when dissolved in water. HCl is an example of a strong acid.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

What is a base?

A

Bases dissolve in water to form Alkalis.They dissociate when dissolved in water to release OH- ions. Therefore can be called proton acceptors

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

What Common bases do you need to know?

A

-Potassium Hydroxide (KOH)
-Sodium Hydroxide (NaOH)
-Ammonia (NH3)

28
Q

What makes a strong base?

A

When dissolved in water it fully dissociates to release OH- ions. NaOH is a strong base.

29
Q

How to form salts by neutralisation?

A

When hydrogen ions in an acid are replaced by metal ions from a base, a salt is formed.

30
Q

Equation for reaction of acid with a metal oxide:

A

Metal oxide + acid –> salt + water

31
Q

Equation for reaction of acid with a metal carbonate:

A

Metal carbonate + acid –> salt + water + carbon dioxide

32
Q

Define shells

A

Energy levels in which electrons orbit the nucleus

33
Q

Which shell has the highest energy?

A

The shell furthest from the nucleus has the highest energy level

34
Q

How many electrons are in each shell?

A

-1st = 2
-2nd = 8
-3rd = 18
-4th = 32

35
Q

What are orbitals?

A

Orbitals are regions within a shell where electrons may be found. Each orbital contains a maximum of 2 electrons.

36
Q

What are the first 4 types of orbitals?

A

S-, P-, D- and F- orbitals

37
Q

What is the energy order of orbitals?

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f e.t.c.

38
Q

What are the orbital shapes?

A
  • S-orbital = spherical
  • P-orbital = dumb-bell (Each p sub-shell has 3 p-orbitals)
39
Q

What are electron configurations?

A

Electron configurations are the arrangement of electrons in atoms or ions e.g.
- Boron, 1s2, 2s2, 2p1
- Nitrogen, 1s2, 2s2, 2p2

40
Q

Define Ionic Bonding

A

Ionic bonds are electrostatic attractions between positive and negative ions

41
Q

Explain what an Ionic compound is

A

An ionic compound has a giant structure and is arranged in a regular giant ionic lattice. The large number of strong electrostatic attractions between ions holds the structure together and gives these compounds high melting points.

42
Q

Define covalent bonding

A

A covalent bond is an electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.

43
Q

What is it called when a covalent bond is formed from an electron pair from one atom?

A

A dative covalent bond

44
Q

Explain electron pair repulsion theory

A

This states that bonding pairs of electrons will repel each other equally in a way that maximises the distance between them

45
Q

Shapes of molecules (without lone pairs):

A
  • 1 electron pairs = linear e.g H2
  • 2 electron pairs = linear e.g CO2, bond angle = 180
  • 3 electron pairs = trigonal planar e.g BF3, bond angle = 120
  • 4 electron pairs = tetrahedral e.g CH4, bond angle = 109.5
  • 5 electron pairs = trigonal bipyramidal e.g PCl5, bond angle = 90 and 120
  • 6 electron pairs = octahedral e.g SF6, bond angle = 90
46
Q

Why do molecules with lone pairs have greater repulsion?

A

An orbital containing a lone pair has greater electron density as a result it has a greater ability to repel than a bond pair

47
Q

How does each bond pair affect the bond angle?

A

Each lone pair reduces a bond angle by approximately 2.5 degrees

48
Q

Shape of molecules (with lone pairs):

A
  • 3 bonding pairs and 1 lone pair = tetrahedral e.g NH3, bond angle = 107 (109.5 - 2.5)
  • 2 bonding pairs and 2 lone pairs = non-linear e.g H2O, bond angle = 104.5 (109.5 - 2 x 2.5)
  • 4 bonding pairs and 1 lone pair = trigonal pyramidal
49
Q

Define electronegativity

A

The tendency for an atom of a given chemical element to attract shared electrons (or electron density) when forming a chemical bond.

50
Q

Explain patterns in electron negativity

A

Electronegativity can be measured using the Pauling scale, which states that electronegativity increases down the group and then across the period with F being the most electronegative.

51
Q

Explain non-polar bonds

A

When both atoms in a bond have identical electronegativity the bond pair of electrons are distributed evenly between the two atoms. An example of this is the bonding between 2 hydrogen atoms another example is the bonding between a carbon and hydrogen atom

52
Q

Explain polar bonds

A

When one atom has a greater electronegativity than the other the bond pair of electrons are distributed unevenly as one atom has a greater attraction. The atom with the higher electronegativity gains a partial negative charge (δ−) subsequently the atom with the lower electronegativity gains a partial positive charge (δ+). The uneven distribution of electrons across the molecule subsequently means the molecule is described as having a dipole. An example of this is the bond between H and Cl within HCl of which the Cl atom has the greater electronegativity

53
Q

Can a molecule contain atoms with different electronegativities and not have a dipole?

A

Yes for example, CCl4 has 4 polar bonds between C and Cl in which carbon is more electronegative, however as these bonds are arranged symmetrically with the molecule there is no overall dipole

54
Q

Explain permanent dipole-dipole interactions

A

Polar molecules such as HCl have permanent dipoles due to the differing electronegativity of the atoms within it due to this the oppositely charged ends of these molecules are attracted to each other. This weak attractive force is known as a permanent dipole-dipole interaction

55
Q

Explain London forces (induced dipole-dipole interactions)

A

Electrons are randomly moving within the shells of a molecule, this can cause the electron density to be spread unevenly causing an instantaneous dipole, if another molecule approaches a weak attraction forms between the two molecules

56
Q

What are Intermolecular forces?

A

Intermolecular forces are seen as permanent dipole-dipole interactions and London forces are the forces of attraction between molecules. These can explain boiling points and solubility

57
Q

How does position in the group affect London forces?

A

As you go down the group the number of electrons per atom increases, which leads to increased strength London forces so more energy has to be applied to overcome these intermolecular forces, hence leading to a higher boiling point for example He has 2 electrons and a boiling point of -269 (degrees celsius) whilst Rn has 86 electrons and a boiling point of -62 (degrees celsius).

58
Q

Order the type of bonds in increasing strength:

A

-London forces
-Permanent dipole-dipole interactions
-Covalent bonding
-Ionic bonding

59
Q

Explain Hydrogen bonding

A

The attraction between an electron-deficient hydrogen atom δ+ and a lone pair of oxygen, fluorine or nitrogen atoms. O, N or F are the only atoms that can form Hydrogen bonds as they are small and highly electronegative. Water molecules can form Hydrogen bonds between each other. Hydrogen bonds are especially strong intermolecular forces.

60
Q

How does Hydrogen Bonding affect the boiling point of Water?

A

Water’s predicted boiling point is -100 degrees Celsius when just looking at London forces as you follow the trend in boiling points of G6 Hydrides however, due to water molecules’ ability to form hydrogen bonds between each other its actual boiling point is +100 degrees celsius

61
Q

Why is ice less dense than water?

A

Hydrogen bonding between water molecules forms a lattice when frozen as ice leading to a rigid structure which means the molecules aren’t packed as close together leading to a lower density

62
Q

What are Simple Molecular Lattices?

A

Many covalently bonded molecules form simple molecular lattices. These structures have strong covalent bonds between the atoms within the molecule. However, the forces between the molecules are weak intermolecular forces, for instance, London forces.

63
Q

What are the properties of simple molecular structures?

A

-Have low melting and boiling points as the intermolecular forces between molecules are weak so a relatively small amount of energy is needed to change state
-Don’t conduct electricity as the structures have no free electrons or free ions
-Dissolve in solvents with intermolecular forces of similar strength

64
Q

What is the strength of Covalent bonds?

A

Covalent bonds are all strong but not all have the same strength, even the same bond may have a different strength depending on the environment e.g oxygen-hydrogen bond in water is different to that in ethanol

65
Q

Explain average bond enthalpy

A

The average amount of energy needed to break one mole of a bond, this can be used to compare the strength of covalent bonds as the bigger the value the stronger the bond

Chemistry Year 1 (4 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions