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Semester 1 Revision > Chemistry Clinical Sciences 7-10 > Flashcards

Chemistry Clinical Sciences 7-10 Flashcards

1
Q

What is a Chemical Bond?

A

A CHEMICAL BOND is the attractive force that holds two or more atoms together in a molecule.

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2
Q

What does bonding involve?

A

The valence electrons.

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3
Q

What do atoms combine? And how can they combine to do this?

A

Atoms combine to gain STABILITY. Electrons in atoms are arranged in ORBITALS so the atom is the most stable it can be.

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4
Q

In bonding what does the rearrangement of electrons allow both ions and molecules to do? Why?

A

Allows molecules/ions to have an even lower energy because when electrons transfer at least one of the atoms will fill less sub levels with a lower energy

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5
Q

What is the Octet Rule?

A

Elements GAIN/LOSE/SHARE electrons to have 8 ELECTRONS in their highest energy level so they have the same electronic configuration as a noble gas.

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6
Q

How many electrons do transition metal have in their highest energy level? What rule does this obey?

A

18 electrons- 18 electron rule.

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7
Q

What is

electronegativity?

A

The ability of an atom to ATTRACT the electrons towards itself in a CHEMICAL BOND.

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8
Q

What does electronegativity depend on?

A

Number of PROTONS in the nucleus.
SIZE of atom.
Amount of SCREENING by inner electron, e.g. amount of SHIELDING between the nucleus and electrons in the highest energy level.

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9
Q

What is the most electronegative element?

A

Fluorine

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10
Q

What is electronegativity used to do in both covalent and ionic bonding?

A

Electronegativity is used in COVALENT BONDING to predict POLARITY of BONDS and in IONIC BONDING to see easy an ATOM will TRANSFER ELECTRONS.

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11
Q

What bonding does low electronegativity + low electronegativity show?

A

Metallic Bonding

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12
Q

What bonding does high electronegativity + high electronegativity show?

A

Covalent Bonding

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13
Q

What bonding does a large difference in electronegativity show?

A

Ionic Bonding

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14
Q

When are ionic bonds formed?

A

IONIC BONDS are formed when one or more electrons is fully TRANSFERRED between atoms. Ionic bonding is between METALS and NON- METALS.

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15
Q

In ionic bonding what do metals do?

A

Metal atoms LOSE electrons to form positive ions (CATIONS).

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16
Q

In ionic bonding what do non-metals do?

A

Non- metal atoms GAIN electrons to form negative ions (ANIONS).

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17
Q

How does an ionic bond form?

Formation of an ionic bond causes a decrease in what? Compared to what?

A

The positive and negative ions ATTRACT each other, electrostatic interaction, to give an ionic bond.

DECREASE in energy compared to the two isolated atoms.

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18
Q

Properties of an ionic bond:

A

Formed due to ELECTROSTATIC ATTRACTION between OPPOSITELY CHARGED IONS.
NON- DIRECTIONAL which means the strength of interaction between two ions depends upon distance not the direction.
Ionic compounds DISSOLVE and
disassociate in POLAR SOLVENTS.
Bonds break when substance is MELTED.
Ionic compounds have NO OVERALL CHARGE.

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19
Q

How does an ionic compound dissolve and disassociate in polar solvents?

A

The bond is broken, and separated ions are surrounded by solvent molecules and interact with solute ions/ molecules.
The bond is broken between oppositely charged ions since the negative ion is attracted to the slightly positive water and the positive ion is attracted to the slightly negative oxygen. If the attraction between the ion and the H/O is strong enough the bond will break.

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20
Q

Features of an ionic bond:

A

METAL and NON-METAL
DISSOLVES easily in WATER and other polar solvents.
CONDUCTS ELECTRICITY when MOLTEN and in SOLUTION.
Form CRYSTALLINE SOLIDS with HIGH melting temperatures.

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21
Q

How is an ionic lattice formed? What pattern does it repeat?

A

OPPOSITELY CHARGE IONS attracted to each other which results in a GIANT IONIC LATTICE which is a repeating 3d pattern of ANIONS and CATIONS in a SOLID state.

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22
Q

What is Lattice Energy defined as?

A

The LATTICE ENERGY is defined as the enthalpy of formation (energy change) of forming one mole of the ionic solid from its gaseous ions- measuring bond strength.

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23
Q

Factors favouring an ionic bond:

Forming a Cation and Forming an Anion. Explain each answer.

A

Forming a Cation- Low ionisation energy of a metal (Lose electrons easily).
Forming an Anion-
High electron affinity/ electronegativity of a non- metal (Gain electrons easily).
Small size of non- metal (Stronger attraction to electrons because of less shielding).
Low charge of anion (Greater attraction to electrons so easier to add electrons).

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24
Q

Other factors which favour an ionic bond forming:

A

Large electronegativity difference between combining atoms. High lattice energy of ionic compound (More energy required to break the bonds so there is a strong attraction between ions).

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25
Q

What is polarisation?

A

The DISORTION of the ELECTRON CLOUD in an atom.

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26
Q

What is polarising power?

A

The ability to distort electron cloud of another ion.

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27
Q

What atoms have a higher polarising power?

A

Smaller the ion and higher charge.

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28
Q

Who has more polarising power cations or anions? Why?

A

Cations are smaller than anions so have a more polarising power whereas anions tend to be polarised.

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29
Q

What atoms have higher polarisability?

A

The larger an atom and the more electrons the atom has the higher the POLARISABILITY.

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30
Q

What shape is an ideal ionic compound?

A

Completely SEPERATE and are SPHERICAL ions.

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31
Q

What are the shape of ions and what character do they have if they are polarised?

A

NON- SPHERICAL ions and COVALENT CHARACTER in bonds

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32
Q

When will an ionic compound have some covalent character?

A

CATION is SMALL and/ or has a HIGH CHARGE which is highly polarising.
ANION is LARGE and/ or has a HIGH CHARGE which is highly polarising.

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33
Q

What are Polyatomic ions?

A

Ions which are made up of MORE THAN ONE ION.

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34
Q

What is the bonding in polyatomic ions?

A

Bonding between atoms in polyatomic ions is COVALENT.

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35
Q

What is a key feature of a polyatomic ion which allows them to get involved in ionic bonding?

A

Number of electrons in a polyatomic ion means the group has an OVERALL CHARGE so can get involved in ionic bonding.

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36
Q

How do you name an ionic compound?

A
  1. First word is cation and usually the parent name.
  2. Oxidation state of metal may sometimes be given in name.
  3. The second word is the ANION. If the ion is polyatomic then the name of the polyatomic ion is used. If the second ion is an element then the end of the element is change to -ide.
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37
Q

How are covalent bonds important?

A

Used in ORGANIC COMPOUNDS which are the building blocks of life.
Determines the SHAPE of molecules which determines interactions between proteins and drugs.
Causes permanent ENZYME INHIBITION.
Determines POLARITY which determines INTERMOLECULAR FORCES so the molecules properties, solubility and receptor interactions.

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38
Q

What are covalent bonds?

A

ELECTRON PAIRS that are SHARED by both participating atoms.

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39
Q

Why do atoms share 1,2 or 3 electron pairs?

A

Want to have the same electronic configuration as a noble gas to be stable.

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40
Q

Features of Covalent bonding:

A

Covalent bonds are a DIRECTIONAL BOND which means a distortion of the electrons are towards an atom.

Covalent bonding is between NON- METALS.

Covalent bonds are formed through SHARED electrons and ELECTROSTATIC ATTRACTION between oppositely charged nuclei and electrons. This attraction forms a covalent bond.

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41
Q

What happens to the number of orbitals when atoms interact to form molecules?

What orbitals do electrons fill from first?

A

The number of orbitals stays the same for example 2 x 1s orbitals = 2 molecular orbitals.

The orbital with the lowest energy.

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42
Q

What combinations are molecular orbitals?

A

LINEAR (on the same plain).

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43
Q

In a correlation diagram if there are 2 anti bonding orbitals full and 3 bonding orbitals full, what happens?

A

2 antibonding orbitals (4 electrons) are full so they cancel out two bonding orbitals (4 electrons) so F2 makes a single bond.

44
Q

How many electrons in:

i) a single bond.
ii) a double bond.
iii) a triple bond.

A

i) 1 pair of electrons- 2 electrons
ii) 2 pairs of electrons- 4 electrons.
iii) 3 pair of electrons- 6 electrons.

45
Q

What do Lewis structures show?

A

Lewis structures are used to show how many bonds an element can make.

46
Q

Describe how you would work out a lewis structure of an element.

A
  1. Valent electrons in each of the elements.
  2. Use pairs of electrons to form a single bond between atoms.
  3. Arrange remaining electrons around atoms to satisfy “octet rule.”
  4. If you run out of electrons use multiple bonds.
47
Q

Describe and explain the factors favouring covalent bonding.

A

High ionisation energies (difficult to lose electrons).
Equal electron affinities. (Atoms attract electrons equally so no polarity so no ionic character).
Equal electronegativities (atom attract the electron pair in covalent bond).
High nuclear charge (able to attract electrons).
Small atomic size (can attract electrons more strongly).
Number of valence electrons (eg. carbon difficult to lose/ gain 4 electrons).

48
Q

Why do elements which favour covalent bonding have high electronegativity values but little difference between these values?

A

Elements should not ionize and there should be equal attraction to electrons between elements.

49
Q

When is a bond NOT polar? Why?

In a NON-POLAR bond where are the electrons? How do we know this?

A

Bonds are non- polar if the two atoms have EQUAL or very similar electronegativity values. Electrons are SHARED EVENLY between the two atoms and tend to be located halfway between the two atoms.
Electrons in the MIDDLE of the bond because the bond has NO DIPOLE, so there is NO SEPERATION of charges within the molecule.

50
Q

When is a bond POLAR? Why?

Polar bonds have a dipole, which leads to?

A

Bonds are polar if the two atoms have small difference in electronegativity values. There is a separation of charge between one end and the other, with greater electron density around the more electronegative atom.
There is a separation of charges.

51
Q

What do polar bonds affect?

A

Reactivity of the bond and the types of intramolecular forces between molecules.

52
Q

Summary of Polar and Non- Polar Covalent Bonds:
No electronegativity different..
Small electronegativity difference…
Large electronegativity difference..

A

No electronegativity different between two atoms leads to a pure non- polar covalent bond.
A small electronegativity difference leads to a polar covalent bond.
A large electronegativity different leads to an ionic bond.

53
Q 
Comparison between Properties Covalent/ Ionic:
State
Melting/ Boiling Point
Conductivity 
Solubility
A

State: C- Gases, liquids or solids + I- Crystalline solids.
M and B: C- Depends on size and intermolecular bonding + I- High.
Conductivity: C- Mostly poor + I- Good when molten.
Solubility: C- Depends on intermolecular bonding + I- Most are soluble in water but not in no polar liquids.

54
Q

Why are Group IV elements good at firing GIANT structures?

A

Form up to 4 Strong bonds so can contain many atoms like an ionic lattice.

55
Q

What are the general properties of giant structures? For example.

A

Very strong, hard, non- conductive insoluble in all solvents.
Graphite can conduct in one direction due to its structure.
Silicon is a semi-conductor.

56
Q

Naming Covalent compounds:

A

Second element has an -ide ending.
Each element has a prefix indicating the number of atoms (not valence).
Eg: For first element drop mono- CO2 = carbon dioxide
The first vowel is often dropped to avoid the combination of “ao” or “oo”.
CO = carbon monoxide (monooxide).

57
Q

Naming Covalent Compounds- list the Prefixes:

A 
1- Mono 
2- Di
3- Tri
4- Tetra
5- Penta
6- Hexa
7- Hepta
8- Octa
9- Nona
10- Deca
58
Q

What is a dative covalent bond?

A

DATIVE COVALENT BOND is a COVALENT BOND that is formed by sharing of two electrons but both of the electrons are from ONE atom.

59
Q

List some of examples when a dative co-ordinate bond is used:

A 
Transition Metal Complex Ions
Ammonium Ion 
Hydroxonium ion
Carbon Monoxide 
Nitric Acid
60
Q

What is the atom which supplies the pair of electrons.

A

The Donor

61
Q

What is the other atom called which is involved in the dative covalent bond?

A

The Acceptor

62
Q

Describe 2 of the requirements for the formation of a dative covalent bond.

A
  1. The atom acting as the donor must have a LONE PAIR of electrons.
  2. The atom acting as an acceptor should have a VACANT orbital to accept the electron pair donated by the donor.
63
Q

What is a lone pair?

A

A LONE PAIR is a pair of electrons in the valence shell of the atom that are NOT INVOLVED in BONDING.

64
Q

Examples of molecules where a dative covalent bond is present.

A

Ammonium ion and Aluminium Complex Ion.

65
Q

Why is dative covalent bonding important in biology?

A 
Function of some biological molecules depends on binding a metal ion-containing cofactor – porphyrins, corrins (via dative bonding). 
For example:
Haemoglobin 
Myoglobin
Cytochromes P450
Vitamin B12
Chlorophyll
Photodynamic therapy
66
Q

Why is dative covalent bonding important in medicine?

A

Function of some drugs depends on their ability to act as donors in dative covalent bonds, for example, chelation therapy & act as acceptors for biological donors.
For example:
Cis platin- Anticancer agent.
Dimercaprol- Treatment of mercury, lead and arsenic poisoning.
D- penicillinamine- Treatment of copper and arsenic poisoning.
Desferoxamine- Treatment of iron overdoes.

67
Q

In simple covalent bonded molecules ions how can the these shapes be predicted. When will this theory only work.

A

Using Valence Shell Electron Pair Repulsion theory.

One central atom and surrounding atoms are approximately the same size.

68
Q

Which has more repulsion bonding pairs/ lone pairs?

A

Bonding Pair - Bonding Pair < Lone pair- Bonding pair < Lone Pair - Lone Pair.

69
Q

Using VSEPR describe how you work out the shape of molecules?

A
  1. Find the number of valence electrons in the central atom.
  2. For ions, add 1 valence electron for anions and remove one for cations.
  3. Work out the number of bonding pairs and lone pairs by drawing the Lewis structure of the molecule.
70
Q

What is electronic geometry?

A

Electronic Geometry is one of the five basic types and SAME for both molecules with lone pairs and no lone pairs.

71
Q

What is molecular geometry?

A

Molecular geometry is DIFFERENT from Electronic Geometry if you have LONE PAIRS- name based on arrangement of atoms.

72
Q

How do you predict the polarity of molecules?

A
  1. If the bonds are NOT polar then the molecule WILL NOT be polar.
  2. Determine the shape of the molecule.
  3. If polar bonds point in the same direction then the molecule is polar but if the point in opposite directions the opposing dipoles cancel so polar is NOT polar.
73
Q

How do you know the direction of the bond?

A

Draw an arrow from the slightly positive atom to the slightly negative atom

74
Q

Describe the polarity of ionic, covalent and dative covalent bonds.

A

Ionic- Polar Bond
Covalent- Can be polar or polarised.
Dative Covalent- Often polar (Old name “Semi- polar bond).

75
Q

What is metallic bonding?

A

ELECTROSTATIC ATTRACTION between METAL IONS and mobile DELOCALISED ELCTRONS. Delocalised electrons mean the electrons are not associated with a particular atom.

76
Q

What are the requirements for metallic bonding?

A

Metals should have LOW IONISATION ENERGIES because metals readily give up their electrons to form positive ions.
Metals should have VACANT VALENCE ORBITALS so electrons move freely between the overlapping orbitals.

77
Q

What are the characteristics of metals?

A

STRONG- Depending of NUMBER OF ELECTRONS because more electrons stronger attraction to positive ions & PACKING OF CATIONS more cations stronger attraction to electrons.
GOOD CONDUCTORS OF HEAT AND ELCTRICITY in solid and liquid states- Electrons can MOVE so PASS on charge/ heat throughout the solid.
MALLEABLE and DUCTILE- Positive ions in LAYERS and layers can SLIDE over each other.
LUSTROUS- Shiny- Bands of electrons REFLECT light.

78
Q

What are INTRAmolecular forces?

A

INTRAMOLECULAR FORCES hold atoms together WITHIN a molecule. For example, covalent and ionic bonds.

79
Q

What are INTERmolecular forces?

A

INTERMOLECULAR FORCES are attractive forces BETWEEN molecules. For example, hydrogen bonding.

80
Q

Which forces intermolecular or intramolecular are weaker?

A

Intermolecular forces are a lot WEAKER than intramolecular forces.

81
Q

What cell interactions are intermolecular forces involved in?

A

Antibody and Antigen complexes
Enzyme- Substrate Complexes
Protein to protein.

82
Q

What attractive/ repulsive forces operate between electrons/ protons when atoms/ molecules approach each other?

A

Electron/ Electron = REPEL
Proton/ Proton = REPEL
Electron/ Proton = ATTRACT

83
Q

List the types of intermolecular forces from strongest to weakest.

A

Ion- Dipole Forces
Hydrogen Boning
Dipole- Dipole Forces
Dispersion Forces- Van der Waals/ London Dispersion forces.

84
Q

Why is there an instantaneous temporary dipole produced in an atom?

A

Electrons are constantly moving in their orbitals. At any one time more electrons may be on one side of the atom/ molecule which produces an instantaneous temporary dipole.

85
Q

If an instantaneous temporary dipole is produced from one molecule what does this create?

What does this result in?

A

Since electrons have moved in the first atom/ molecule, the one end of the first atom/ molecule is slightly positive and the other end slightly negative. Therefore another TEMPORARY DIPOLE is INDUCED in another molecule/ atom because like charges repel and opposite charges attract.

Results of the temporary dipole are WEAK INTERMOLECULAR FORCES between molecules/ atoms.

86
Q

Give another name for London Dispersion forces.

Between what molecules do London Dispersion forces occur?

A

Van de Waals forces.

All molecules.

87
Q

What molecules have the strongest intermolecular forces? Why?

A

LARGER molecules have STRONGER intermolecule forces because they have more electrons so stronger attraction and repulsion forces.

88
Q

What is the significance of London Dispersion Forces?

A

Present in all molecules but most important in NON POLAR molecules because only forces between these molecules.
Determine BOILING POINT of noble gases, halogens and alkanes.
Determine HARDNESS of certain solids.
Determines GAS BEHAVIOUR at a lower temperature.

89
Q

What are intermolecular forces important in stabilising? Give an example.

A

Lipid bilayers, for example the cell membrane and PROTEIN STRUCTURE and PROTEIN PROTEIN recognition.

90
Q

What interactions relate to PERMANENT molecule dipoles?

A

DIPOLE- DIPOLE INTERACTIONS

91
Q

How does a molecule have dipole- dipole interactions?

A

Bonds have a difference in electronegativity.

The molecular dipole require a NON-UNIFORM CHARGE DISTRIBUTION in 3-D.

92
Q

Explain what dipole- dipole interactions are.

What do these forces result in for a molecule?

A

Dipole- dipole interactions is ELECTROSTATIC ATTRACTION between PERMANENT DIPOLES and act in addition to London interactions.

These forces result in a HIGHER BOILING POINT than expected from the mass of the molecule.

93
Q

What molecules are dipole- dipole interactions significant in?

A

Carbonyls and Protein Folding

94
Q

What are ion- dipole forces?

What does the strength of ion- dipole forces depend on?

A

ION- DIPOLE FORCES is ELECTROSTATIC ATTRACTION between an ION and an UNCHARGED POLAR MOLECULE. Its STRENGTH depends on CHARGE of the ion and MAGNITUDE (size) of the dipole.

95
Q

Where are ion- dipole forces most commonly found? Give an example.

A

In solution, for example KCl dissolving in water.

96
Q

What molecules do hydrogen bonding occur between? Why?

A

Hydrogen bonding occurs between an ELECTRONEGATIVE ATOM, NITROGEN, OXYGEN or FLURORINE because they are small in size, have a high charge density and lone pairs of electrons in small orbitals.

97
Q

Which molecules are known as the Acceptor? What molecule is known as the Donor?

What type of bond is between atoms?

A 
Acceptor= Nitrogen, Oxygen or Flurorine. 
Donor= Hydrogen.

Polar Bond.

98
Q

What are the properties of hydrogen bonds?

A

WEAK and easily broken in comparison to covalent bonds.
DIRECTIONAL
ADDITIVE

99
Q

What is the significance of hydrogen bonding?

A

Large numbers of hydrogen bonds account for the PHYSCIAL STATE of WATER.
Strong enough for STABILITY of MACROMOLECULES.
Needed in DNA double helix BASE PAIRING.
Needed in Protein secondary structure in ALPHA HELICES and BETA SHEETS.
Needed in ENZYME- SUBSTRATE and ANTIBODY- ANTIGEN INTERACTIONS.
Important for the SOLUBITILTY of COMPOUNDS in WATER.

100
Q

If a hydrogen bond is present between two molecules instead of elements within a molecule, what does this mean for the first substance?

A

The substance will have a higher melting point than a substance with hydrogen bonds between elements in a molecule/ atom.

101
Q

Why does water have an extensive network of hydrogen bonding?

A

On each oxygen atom there is 2 LONE PAIRS and each water molecule has 2 HYDROGEN ATOMS.

102
Q

How many hydrogen bonds can water form? What shape does this make water?

A

4 hydrogen bonds= Tetrahedral Formation.

103
Q

What structure is ice similar to?

A

Diamond- Crystal Structure.

104
Q

What is the difference between liquid water and ice?

A

Molecules in water are further apart than in liquid water.

105
Q

Why is Ice less dense that liquid water?

A

In water hydrogen bonds break and reform but in ice hydrogen bonds in a fixed position which is the optimum position from hydrogen bonds.

106
Q

Examples of hydrogen bonding in biology.

A

Matching number of hydrogen bond donors and acceptors allows ‘correct’ pairing of DNA bases.
Hydrogen bonding is needed in alpha helices and beta sheet structures.
Hydrogen bonding used in Cellulose which is a polymer of glucose. Sugar chains are held rigidly together by multiple hydrogen bonds within and between chats.

Semester 1 Revision (9 decks)

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