BMS165 Topic 2: Chemistry

Elements

• Matter is defined as anything that has mass and takes up space (it has volume).

• All matter is composed of atoms, and atoms are indivisible and indestructible.

• An element is a substance that is made up of only one kind of atom.

• Simplest form of matter to have unique chemical properties

• Each element defined by its Atomic number based on the number of protons in its nucleus

• The Periodic table (Elements arranged by atomic number & represented by one- or two-letter symbols)

Major Elements

• 118 elements have been identified – 94 exist naturally on earth

• 24 elements have biological roles

• 6 elements = 98.5% of body weight

  1. Oxygen (O)

  2. Carbon (C)

  3. Hydrogen (H)

  4. Nitrogen (N)

  5. Calcium (Ca)

  6. Phosphorus (P)

• Trace elements - present in minute amounts, but play vital role

Elements make up minerals

• Inorganic elements extracted from soil by plants and passed up food chain to humans

• Constitute about 4% of body weight

• Ca and P make up about 3%

• Remaining 1% is mainly Cl, Mg, K, Na, and S

• Important for body structure

  1. Ca crystals in teeth, bones, etc.

  2. Important for enzyme function

  3. Electrolytes are mineral salts needed for nerve and muscle function

Atoms

• Atoms cannot be cut or individualised

• Smallest unit of matter that retain the properties of elements

• Generally, refers to one individual particle of the smallest unit of matter.

• Atoms are very small: a ten-billionth of a meter.

• Elements and atoms cannot be broken down by chemical means into simpler substances.

• The radius of an atom is about 1 to 2 x 10-8 cm

Atomic Structure

• The nucleus lies at the centre of an atom.

• It is made up of a number of particles that are smaller than the atom (i.e. sub-atomic particles)

• We will only deal with two of these:

  1. Protons (positive charge)

  2. Neutrons (no charge)

Protons

• Protons are POSITIVELY charged particles

• Protons give atoms their identity

• The number of protons determines what element the atom it is

• The number of protons present in an atom is known as the atomic number

• Every time you add a proton, the atomic number goes up by one = different element (with different chemical properties).

• The atomic number = the number of protons!

Neutrons

• Neutrons are a sub-atomic particle in the nucleus

• Neutrons possess NO CHARGE

• Neutrons help keep the nucleus together

• Acts as a buffer between the positive protons within the nucleus

Neutrons + Protons equals Atomic Mass

Electrons

• Electrons are a sub-atomic particle outside of the nucleus

• Have a lot less mass than protons and neutrons

• The number of electrons is always equal to the number of protons found in an element

• Remember that the protons are positively charged and the electrons are negatively charged

• Positive charge and negative charge are attracted to each other

• It is this electromagnetic attraction between positive and negative charge that holds the electrons “close” to the nucleus (relative to the size of these particles, the distance is actually quite large

Electron Shell

• One of the main chemical properties of elements is how they form stable “bonds” (or not) with other elements

• How atoms bind to each other is largely dependent on the behaviour of electrons

• Electrons are only found in certain places around the nucleus, these places are called electron shells

• Not all the electrons of an atom can exist in the same electron shell

• The more electrons there are in an atom the further away from the nucleus they have to go into different shells (think onion/tree rings).

• The first electron shell only holds up to 2 electrons.

• Once first shell is full electrons move further away from the nucleus to the second electron shell.

• The second electron shell holds up to 8 electrons.

• Once full electrons need to move to the third shell

• The third electron shell also holds up to 8 electrons.

Valence Shells

• Valence Shells: Refers to the outermost electron shell that contains electrons in an atom.

• Example 1 Neon

  • valence shell is the second electron shell.

• Example 2 Chlorine

  • Valence shell is the third electron shell

• This shell determines the binding properties of the atom and therefore how it reacts

Octet Rules

• Atoms are most stable when their valence shell is full

• Elements such as these have “full” valence shells

  • Helium (2)

  • Neon (10)

  • Argon (18)

• These elements have similar chemical properties so they are grouped together in the same column on the periodic table

• Noble gasses are stable = have full valence shell.

• No other elemental atoms have full valence shell

• Other elements achieve stability by

  • generating a “full” valence shell configuration similar to the Noble gases by “sharing” electrons with each other in chemical bonds.

• This tendency of atoms to prefer to have eight electrons in their second and third electron shells is known as the octet rule (2:8:8).

• This is why we have molecules: atoms interacting with each other to share or swap electrons and get every atom to have a full valence shell.

Elements, Molecules and Chemical Bonds

• Element

  • The simplest substance that cannot be broken down any further.

• Molecule

  • two or more atoms that are chemically bound.

• Compound

  • A compound is a molecule that is made up of 2 or more elements.

• Molecular formula

  • gives the kind and number of atoms of each element present in a molecular compound CO2

• Structural formula

  • Identifies location of the chemical bonds between atoms of a molecule

Molecules and Chemical Bonds

There are 3 types of Chemical Bonds that are important for the human body

1. Ionic bond

   • Exchanging electrons

   • Strong

2. Covalent bond

   • Sharing electrons

   • Strong

3. Hydrogen bond

   • Attractions of slightly positive H atoms for other molecules

   • Weak

Isotopes and Radioactivity

• Isotopes are varieties of an element that differ only in the number of neutrons

• Isotopes = same number of protons + different number of neutrons

• Extra neutrons increase atomic weight

• Isotopes of an element are chemically similar because they have the same number of valence electrons

• Atomic weight (relative atomic mass) of an element accounts for the fact that an element is a mixture of isotopes

• For example. Carbon has an atomic number of 6 but an atomic weight of 12.011 because of the existence of small amounts of 13C and 14C.

• Hydrogen is the first element in the periodic table and has the atomic number one. Those elements which have same atomic number but different mass number are called isotopes. There are three isotopes of hydrogen namely, protium, deuterium and tritium.

• Radioisotopes

  • Unstable isotopes that decay and give off radiation

  • Every element has at least one radioisotope

• Intense radiation can be ionizing

  • Ejects electrons, destroys molecules, creates free radicals

  • Can cause genetic mutations and cancer

  • Examples: UV radiation, X-rays, alpha particles, beta particles, gamma rays

• Ion

  • Charged particle (atom or molecule) with unequal number of protons and electrons

• Ionisation

  • Transfer of electrons from one atom to another

• Anion

  • Particle that has a net negative charge (due to gain of electrons)

• Cation

  • Particle that has a net positive charge (due to loss of electrons)

  • Ions with opposite charges are attracted to each other

Ions, Electrolytes, and Free Radicals

• Electrolytes

  • Substances that ionise in water and form solutions capable of conducting electric current

• Electrolyte importance

  • Chemical reactivity, osmotic effects, electrical excitability of nerve and muscle

  • Electrolyte balance is one of the most important considerations in patient care (imbalances can lead to coma or cardiac arrest)

• Free Radicals

  • molecular species capable of independent existence that contains an unpaired electron in an atomic orbital

  • There are many types of radicals, but those of most concern in biological systems are derived from oxygen, and known collectively as reactive oxygen species. Oxygen has two unpaired electrons in separate orbitals in its outer shell. This electronic structure makes oxygen especially susceptible to radical formation.

• Oxygen-derived radicals are generated constantly as part of normal aerobic life

• Antioxidant: A substance that reduces damage due to oxygen, such as that caused by free radicals.

Water and Mixtures

• Mixtures​

  • Consist of substances that are physically blended but not chemically combined​

  • Body fluids are complex mixtures of chemicals​

• Water​

  • Most mixtures in our bodies consist of chemicals dissolved or suspended in water​

  • Water is 50–75% of body weight​

  • Depends on age, sex, fat content, and so on

Water

• Polar covalent bonds and a V-shaped molecule give water a set of properties that account for its ability to support life​

  • Solvency: the ability to dissolve other chemicals​

  • Adhesion: clings to a lot of different substances ​

  • Cohesion: water molecules cling to one another​

  • Chemical reactivity: water participates in a lot of reactions and itself ionises into H+ & OH-​

  • Thermal stability: changes temperature relatively slowly

• Hydrophilic​ = water loving

  • Substances that dissolve in water​

  • Hydrophilic molecules are polarised or charged (e.g., sugar)​

• Hydrophobic​ = water hating

  • Substances that do not dissolve in water​

  • Hydrophobic molecules are nonpolar or neutral (e.g., fats)

Mixtures

• A combination of elements or compounds that are physically blended together but NOT bound by chemical bonds

• Mixtures can contain solids, liquids and gases either alone or in combination

• Some examples of liquid mixtures are;​

  • Solutions​

  • Suspensions​

  • Colloids

Type of Mixtures: 1. Solutions

• A solution is a mixture of a solute dissolved into a solvent that remains evenly distributed = is homogenous

• For example, sugar (solute) dissolved into water (solvent).​

• Solvent ​

  • is a liquid (or gas) in which another substance is dissolved

  • In the body this is usually water

• Solute​

  • is a substance that is dissolved into a solvent.​

Types of Mixtures: 2. Suspensions

• A suspension is a heterogeneous mixture in which some of the particles settle out of the mixture upon standing

• The particles in a suspension are far larger than those of a solution

  • gravity is able to pull them down out of the dispersion medium (water)

• The diameter for the dispersed particles in a suspension, such as the sand in the suspension described, is typically at least 1000 times greater than those in a solution. ​

• Unlike a solution, the dispersed particles can be separated from the dispersion medium by filtering

• Suspensions are considered heterogeneous because the different substances in the mixture will not remain uniformly distributed if they are not actively being mixed

Types of Mixtures: 3. Colloids

• A heterogeneous mixture in which the dispersed particles are intermediate in size between those of a solution and a suspension

• The particles are spread evenly throughout the dispersion medium, which can be a solid, liquid, or gas

• Because the dispersed particles of a colloid are not as large as those of a suspension, they do not settle out upon standing

• Colloids are unlike solutions because their dispersed particles are much larger than those of a solution

• Emulsions are a type of colloid.

Strong Acid and Strong Base

• A strong acid is a substance that has a strong tendency to dissociate (separate) completely to release H+

  • HCl (hydrogen chloride)

• A strong base is a substance that has a strong tendency to dissociate (separate) completely to release OH-

  • NaOH (sodium hydroxide)

Weak Acid

• A weak acid is a substance that has a weak tendency to dissociate (separate) to release H+

  • may also reform (i.e. reversible)

  • Do not fully dissociate

  • Significant number of molecules remain in solution = reversibleEg: Carbonic acid (H2CO3)

• A weak acid found within the body. In solution carbonic acid reversibly dissociates (separates) into a hydrogen ion and a bicarbonate ion.

Weak Base

• A weak base is a substance that has a weak tendency to dissociate (separate) to release OH- and may also reform (i.e. reversible)

  • They do not fully dissociate

  • A significant number of molecules remain in solution, therefore can be reversible.

• Eg: Ammonia (NH3)

  • In solution ammonia will produce ammonium ions and hydroxide ions.

pH and Its Significance

• Definition of pH: Stands for "power of hydrogen"; it measures the concentration of hydrogen ions in a solution.

• Neutral pH: A pH of 7 indicates a neutral solution where the concentration of hydrogen ions ([H+]) and hydroxide ions ([OH-]) is equal (1:1 ratio).

• Acidity and Basicity: pH values lower than 7 indicate acidity (higher [H+]), while values above 7 indicate basicity (higher [OH-]).

• Examples of pH Values:

  • pH of 6: [H+] is 10 times greater than [OH-] (10:1 ratio).

  • pH of 5: [H+] is 100 times greater (100:1 ratio).

  • Progressively, as pH decreases from 7 to 1, [H+] increases dramatically (e.g., pH 1 corresponds to 10 million times greater [H+] than neutral).

Blood pH

• Normal Blood pH Range: 7.35 to 7.45 - crucial for physiological balance in the body.

• Importance of Blood pH: Deviations can significantly impact cellular functions and overall body metabolism.

Concentration in Solutions

• Concept of Concentration: Refers to the amount of solute relative to the solution; could be expressed in moles/liter or percentage.

• Measure of Concentration:

• Isotonic Solutions: A solution with a concentration of 0.9% sodium chloride is considered isotonic; it maintains cell integrity upon injection.

  • Preparation: 0.9% means 0.9 g of sodium chloride in 100 mL of distilled water.

Carbohydrates

• Definition: Basic sugars and starches that provide energy (e.g., glucose).

• Composition: Made up of carbon (C), hydrogen (H), and oxygen (O); commonly represented as CHO.

• Types of Carbohydrates:

  • Monosaccharides: Single sugar unit (e.g., glucose).

  • Disaccharides: Two monosaccharides (e.g., sucrose).

  • Polysaccharides: Multiple sugar units (e.g., starch).

• Energy Conversion: Glucose is broken down for energy, resulting in Adenosine Triphosphate (ATP).

Lipids

• Definition and Function: Include fats and oils; essential for energy storage and cellular structures.

• Fatty Acids:

  • Saturated: No double bonds between carbon atoms, fully saturated with hydrogen.

  • Unsaturated: Contains one or more double bonds, creating space for additional hydrogen.

• Triacyglycerols (Triglycerides): Composed of one glycerol and three fatty acids; a key storage form of energy.

• Phospholipids: Important for cell membranes; contain hydrophilic (water-attracting) heads and hydrophobic (water-repelling) tails.

Proteins

• Basic Structure: Composed of amino acids with a general formula of NH2 (amino group) and COOH (carboxyl group).

• Levels of Protein Structure:

  • Primary: Sequence of amino acids.

  • Secondary: Coiling or folding of the amino acid chain (e.g., alpha-helices and beta-sheets).

  • Tertiary: Overall 3D shape of a polypeptide.

  • Quaternary: Multiple polypeptide chains assembling together (e.g., hemoglobin).

• Importance of Proteins: Function in catalysis (enzymes), structure (collagen), transport (hemoglobin), and more.

Nucleic Acids

• Definition: Include DNA and RNA, vital for genetic information transfer.

• Composition: Composed of nucleotides, containing a five-carbon sugar, a phosphate group, and a nitrogenous base.

Chemical Reaction

• In a chemical reaction the bonds that hold atoms together are broken and rearranged to make new substances.

• No atoms or matter are created or lost in chemical reactions

  • the amount of reactant atoms must equal the amount of product atoms.

• Energy is the capacity for doing work/causing change

  • making something move/hotter, breaking apart chemical bonds

• The law of conservation of energy states that energy cannot be created or destroyed (it can only change form)

• No matter what you do, the energy you put in = energy you get out

• Energy comes in many forms, but there are two main categories

• Potential energy

  • includes chemical, electrical, nuclear

• Kinetic energy

  • includes mechanical, therm

Exothermic reactions: Chemical reactions = that result in the net release of energy from breaking chemical bonds

Endothermic reactions: absorb energy in the form of heat and stores it in newly formed chemical bonds.

Chemical Reaction Rates

• Chemical reactions occur when molecules collide with enough force and correct orientation

• Chemical reaction rates increase when:

  • Concentration of reactants increases

  • Temperature rises

  • A catalyst is present:

    • Enzyme catalysts bind to reactants and hold them in orientations that facilitate the reaction

    • Catalysts are not changed by the reaction and can repeat the process frequently

Metabolism: the chemical processes that occur within the cells of a living organism in order to maintain life. This includes breakdown and assembly of molecules. As well as reaction energy output or input.

Catabolism: break down of molecules (e.g. in food) produce energy or small subunits that can subsequently be reassembled into larger molecules by anabolism.

Anabolism: synthesis (building) of complex molecules and energy storage compounds in living organisms from simpler molecules.