Post on 19-Dec-2015
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
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
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
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
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
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
Polarity of Molecules - Examples
Water: H2O
VSEPR diagram:
Polar bonds? YesOverall dipole? YesWater is: POLAR
Polarity of Molecules - Examples
Carbon dioxide: CO2
VSEPR diagram:
Polar bonds? YesOverall dipole? NoCarbon dioxide is: NON-POLAR
Polarity of Molecules - Examples
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
Special 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.
Special Properties 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
Special Properties of water
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.
Special Properties of water
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
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