Chemical Fundamentals. Biology is the study of living things All living matter is ultimately...

BiochemistryChemical Fundamentals

Biology is the study of living things All living matter is ultimately composed of

chemical substances Matter is anything that has mass and takes

up space

The nucleus is made up of protons (p+) and neutrons (no) and is surrounded by rings of orbiting electrons (e-)

The Bohr-Rutherford Model of the Atom

Subatomic particle

Relative mass Relative charge

Proton 1 +1

Electron 0 -1

Neutron 1 0

Also called isotope notation

X = element symbolA = atomic mass = # protons + # neutronsZ = atomic number = # protons

Standard atomic notation

XAZ

What is the difference between these two atoms?

Isotopes

Atoms in which the number of neutrons may differ

12C and 14C are two isotopes of carbon In nature, these isotopes differ in

abundance The relative abundance of isotopes is taken

into account to produce the atomic mass you see on periodic tables

m carbon= 12.011 amu

Isotopes

The nucleus on some isotopes spontaneously break apart or decay.

The matter and energy given off in this decay process causes these isotopes to be radioactive.

This results in the formation of new elements When 14C decays, it becomes 14N. The length of time it takes for a radioactive

substance to decay by half is called the half-life. Radioisotopes can be both useful and

dangerous Radiation can cause mutations in DNA, so need

to be handled with care in order to limit exposure

Radioisotopes

radioisotopes are used in medical imaging

Injected isotopes localize in specific tissues and release radiation outwards

this radiation is detected by special cameras

Radioisotopes

Radioisotopes are also useful in tracing molecules in biochemical pathways (a complex series of reactions in a cell).

Molecules which contain a lot of nitrogen (amino acids for instance), can be ‘tagged’ with a radioisotope of nitrogen

Radioisotopes

Are also useful for finding the absolute age of rocks, fossils, or ancient specimens unearthed by archaeologists or palaeontologists

Radiometric dating relies on the half-life of radioisotopes

While an organism is alive, it is taking in carbon and incorporating it into its tissues – all isotopes in their relative amounts.

When it dies, it stops taking in carbon By measuring the amount of parent isotope vs.

daughter isotope, the half-life can be used to calculate how long it has been since the organism stopped taking in the parent isotope

Radioisotopes

Ions are elements that have gained or lost electrons

Ions are commonly found dissolved in water, such as in the cytoplasm or plasma of the blood

Elements in the same family tend to form the same type of ion (e.g.: Na+, Li+, K+, Rb+)

Some important ions are Ca2+ (used for muscle contraction), Na+ and K+ (nerve and muscle function), Fe2+ and Fe3+(in hemoglobin) and H+ (required for synthesis of ATP)

Ions

Electrons orbit the nucleus of an atom at a great distance compared to the size of the particles

Analogy: If an apple represented the size of an atom’s nucleus and it was placed at the center of the earth’s core, the valence electrons would be orbiting close to the surface of the earth’s crust

The valence electrons therefore are the part of the atoms that interact in chemical reactions to form compounds

Chemical Bonding

Form between a metal and a nonmetal Metal tends to lose electrons which are

transferred to the nonmetal Metals form a cation (+) and nonmetals

form an anion (-)

Ionic Bonds

Formation of NaCl Formation of MgF2

Ionic Bonds

These result in a lattice of ions rather than individual molecules, so we refer to MgF2

and NaCl as formula units, not molecules. Properties of ionic substances:◦Crystalline solids at room temperature◦Hard and brittle◦High melting and boiling points◦Conduct electricity when in liquid form◦Most are soluble in water

Ionic substances

Form between two nonmetals Electrons are shared rather than

transferred Macromolecules and organic molecules are

covalent molecules using covalent bonds, such as lipids, carbohydrates, proteins and nucleic acids.

Covalent Bonds

Covalent Bonds

Formation of H2 Formation of NH3

Covalent Bonds

Formation of O2 – a double bond

Linus Pauling developed the concept of electronegativity (En)

It is a measure of how strongly an atom attracts electrons to itself

Fluorine has the highest En value, and Pauling assigned it an arbitrary value of 4.1

Elements to the left and below fluorine have decreasing En values

Electronegativity

In bond formation, it is useful to look at the electronegativity difference (ΔEn)

When Pauling looked at a range of bonds and their ΔEn values, a pattern was noticed

Bonds with an ΔEn between 1.7 – 4.1 tended to exhibit ionic characteristics

Bonds with an ΔEn below 1.7 tended to exhibit covalent characteristics

HBr En (hydrogen) = 2.1 LiF En (lithium) = 1.0

En (bromine) = 2.8 En (fluorine) = 4.1

ΔEn = 2.8 – 2.1 = 0.7 ΔEn = 4.1 – 1.0 = 3.1

Electronegativity difference

Trends in ΔEn

There are two types of covalent bonds Atoms that have the same En will have an

ΔEn of zero. These atoms will attract the shared

electrons equally, and so the distribution of electrons is uniform

These are nonpolar covalent bonds

Polar and Nonpolar covalent bonds

Covalent bonds that have two different elements will have different En values and so the electron distribution will be non-uniform

These bonds are called polar covalent, since one end of the bond will be slightly electronegative (δ-)since the electrons are attracted more to the atom at that end

Polar and Nonpolar covalent bonds

Valence Shell Electron Pair Repulsion theory (VSEPR) allows us to predict the 3-D shape of a molecule

VSEPR theory states that bond pairs of electrons repel one another, and lone pairs of electrons take up more space than bond pairs

There are four basic shapes which are common in organic molecules:

Linear Bent or V-shaped Tetrahedral Pyramidal

Molecular Shape

Linear

or

Molecular shape

Bent

Molecular shape

Pyramidal

Molecular shape

Tetrahedral

Molecular shape

How do we determine if a molecule is polar or nonpolar?

A polar molecule has an uneven distribution of electrons. This occurs when◦There is at least one polar bond ◦The shape of the molecule is asymmetrical◦Or the shape is symmetrical but the atoms

surrounding a central atom have different En values

Polarity of covalent molecules

Methane: CH4

VSEPR diagram:

Polar bonds? YesOverall dipole? NoMethane is: NON-POLAR

Polarity of Molecules - Examples

Ammonia: NH3

VSEPR diagram:

Polar bonds? YesOverall dipole? YesAmmonia is: POLAR


Water: H2O

VSEPR diagram:

Polar bonds? YesOverall dipole? YesWater is: POLAR


Carbon dioxide: CO2

VSEPR diagram:

Polar bonds? YesOverall dipole? NoCarbon dioxide is: NON-POLAR


The particle theory states that there are forces between particles, and the forces increase as the particles get closer.

These are the intermolecular forces Compared to covalent and ionic bonds, they

are very weak – but when there are many, they add up to a significant force

Collectively they are called van der Waals forces, but there are three different forces.

These forces have an effect on the boiling point and the solubility of substances.

Intermolecular Forces

London dispersion forces (LDF) occur when the protons in one atom or molecule attract the electrons in a neighbouring atom or molecule.

Since all particles have protons and electrons, all substances have LDF

Larger molecules have more protons and electrons, and so have greater London dispersion forces.

London Dispersion Forces

London Dispersion Forces

When comparing the boiling points of hydrocarbons (non-polar molecules), we see that the boiling point increases as the number of carbons increases.

Why is this?

Occurs only in polar molecules that have hydrogen and at least one of the following atoms: N, O or F.

These highly electronegative atoms have lone pairs of electrons which are attracted to the hydrogen atoms in neighbouring molecules.

These hydrogen atoms are essentially a proton

Hydrogen bonding

Polar substances have a slightly electronegative end and a slightly electropositive end.

Dipole-dipole forces occur when oppositely charged poles momentarily attract one another

Dipole-dipole forces

Water is not an organic molecule but is essential for life on this planet

All cells are surrounded inside and out with water – anything that interacts with a cell must first be dissolved in water

Physical properties:◦ colourless and transparent◦ liquid at room temperature◦density = 1.0 g/mL◦m.p. = 0℃ b.p = 100℃

water has LD, D-D forces, and H-bonding

Water

Water has cohesive properties – the high number of intermolecular forces causes water molecules to ‘stick’ together

Examples:◦ surface tension – beading of water◦water striders – too light to break surface tension◦ transpiration in plants – transport in xylem tubes

Special Properties of Water

Water has adhesive properties – it’s polar nature causes it to stick to other substances

Examples:◦ capillary action – water ‘climbs’ up small diameter

tubes, or ‘bleeds’ through the microscopic pores and channels in paper or other porous substances

◦ this is due to the hydrogen bonding interactions between the water and the surface of the tube (either SiO2 or the cellulose tubes of paper)

◦This helps to explain the meniscus inside a tube


Cohesive-Adhesive Properties of Water

Water has outstanding solvent properties Used to be called the ‘universal solvent’, but this

is not a good name, since not everything dissolves in water

The polar nature of water allows any other polar substance or any charged particle to dissolve easily

The δ- will attract the δ+ end of solutes, and this attraction will remain once the solute is dissolved.

The same is true for ionic substances – the cation will be attracted to the δ- end of water, and the anion will be attracted to the δ+ of water.


Water has a high specific heat capacity This is a measure of the amount of heat energy

required to increase the temperature of a 1g of a substance by 1℃.

cwater = 4.18 J/g‧℃

This is high compared to other substances: ccopper = 0.385 J/g‧℃ cair = 1.00 J/g‧℃

cglass = 0.735 J/g‧℃ ciron = 0.450 J/g‧℃

A metal pan absorbs heat energy quickly and loses it quickly. This makes metals useful for cooking.

Water takes more energy to heat up – thus the time it takes to boil water in a pot.

Special Properties of water

Moderation of climate

This property of water also helps to moderate temperature changes in cells

Specific heat capacity of water

Water has a high latent heat of vaporization and fusion.

Latent heat is the energy absorbed or released by a substance during a change of state.

Lf water = 334 J/g Lv water = 2260 J/g


Latent heat

Evaporative cooling relies on Lv of water.

Latent heat of vaporization

Tender fruit farmers take advantage of the latent heat of fusion of ice when there is a chance of frost

On an evening when there is frost in the forecast, they spray water over their fruit, causing ice to form as the temperature drops below 0°C.

How does this help to protect the fruit?

Latent heat of fusion

Water’s density decreases as it changes from liquid to solid.

This is because the distance between molecules in a crystal lattice (as ice) on average further than when in a liquid.


Ionization occurs when 2 water molecules break apart into a hydronium ion H3O+ and a hydroxide ion OH-

2H2O ⇌ H3O+ + OH-

A substance that releases H+ ions in solution is an acid.

A substance that releases OH- ions in solution is a base.

Acids and Bases

In pure water at 25 ℃, [H3O+] = 1.0 x 10-7 mol/L

pH is a measure of the concentration of hydronium ions [H3O+]

pH = -log [H3O+] When an acid is combined with a base, a

neutralization reaction occurs, with a salt and water as products.

HCl + NaOH ⇌ NaCl + H2O

Acids and Bases

Strong acids and bases ionize or dissociate completely when dissolved in water.

A strong acid produces a high number of excess H+ ions which combine with water to produce H3O+ ions.

Weak acids and bases ionize only partly. Only about 1.3% of weak acid’s molecules

(such as acetic acid) will ionize in water, so contribute far fewer H+ to the solution.

Strong vs Weak acids and bases

Hydrochloric acid 100%

HCl → H+ + Cl-

Strong acids

Acetic acid 1.3%

CH3COOH ⇄ CH3COO- + H+

Weak acids

Biological systems tend to dislike significant shifts in pH. Our environment is full of weak acids and bases that can easily shift biological systems from their optimum pH. Buffer systems prevent these shifts.

Example: blood has an ideal blood pH of ~ 7.4 Below pH 6.8 or above pH 7.8, death occurs CO2 and H+ are produced during cell respiration When blood pH rises, carbonic acid dissociates to form

bicarbonate and H+.H2C03 ⇌ HC03

- + H+

When blood pH drops, bicarbonate ions bind H+ to form carbonic acid.

HC03- + H+ ⇌ H2C03

Acid-Base Buffers

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