CHAPTER 2

: THE

CHEMICAL C

ONTEXT

OF

LIFE

MULTIDISCIPLINARY SCIENCE

Biological systems follow the same laws that are dictated by physics and chemistry.

Emergent properties – biological functions are made possible by the interaction between atoms and molecules that make up an organism.

Discuss:

What is the definition of life?

What is required for something to be considered alive?

At what point do we cross the line between nonliving and living at the molecular level of cell biology?

Do you consider a virus to be “alive?”

Bombardier Beetle (chemistry) - https://www.youtube.com/watch?v=Pib9qT-pccIPistol Shrimp (physics) - https://www.youtube.com/watch?v=QXK2G2AzMTU

THE BASICS - IMPORTANT TERMS

Matter – Anything that has mass and occupies volume.

Element – A substance that cannot be broken down into other substances through chemical reactions.

Compound – A substance that consists of two or more different elements in a fixed ratio.

Atom – The smallest unit of an element that maintains the chemical properties of the element.

Molecule – The smallest unit of a compound that maintains the chemical properties of the compound,

LE 2-2

Sodium Chlorine Sodium chloride

THE BASICS – PERIODIC TABLE

ELEMENTS ESSENTIAL TO LIFE

• 92 natural known elements• 25 elements are essential to life• 4 elements make up 96% of

living matter• Carbon (C)• Oxygen (O)• Hydrogen (H)• Nitrogen (N)

• Remaining 4%• Phosphorous (P) • Sulfur (S)• Calcium (Ca)• Potassium (K)

• Less that 0.01% of living matter is made up of trace elements

THE STRUCTURE OF ATOMS

• Subatomic particles• Neutrons - ~1 amu (1.7 x10ˆ-24 grams); no charge• Protons - ~1 amu; +1e (1.602 x 10ˆ-19 coulombs)• Electrons - 5.4858 x 10ˆ-4 amu (or 1/1840 amu); -1e (-

1.602 x 10ˆ-19 coulombs)

• What makes up the atomic nucleus?

• 1 amu=1 dalton

• Atomic Weight - # of Protons + # of Neutrons• Mass of electrons is negligible• AKA Atomic Mass; Mass Number

• Atomic Number - # of Protons

Atomic Number

Atomic Weight

Atomic Number Atomic Weight

LE 2-4

Nucleus

ElectronsCloud of negativecharge (2 electrons)

• Atoms of the same element that have different atomic weights

• Always have the same number of protons

• Differ in the number of neutrons

• Atomic Weight = Average Atomic Weight of all isotopes

• Some isotopes are stable

• Most isotopes are unstable (radioactive)

ISOTOPES

ISOTOPES – APPLICATIONS IN LIFE SCIENCES• Dating techniques for biological

materials and fossils

• Tracers used to follow metabolic processes (Rate of DNA synthesis)

• Medical Diagnostics• PET Scans

• Risk associated with using radioactive isotopes• Severity depends on type and amount

of radiation

ISOTOPES - PRACTICE

How many protons, neutrons, and electrons are present in the following isotopes?

• Technetium-99

• Flourine-18

• Iodine-131

• Cobalt-60

What is the atomic mass of an element, given the following isotopes?

Oxygen-16, oxygen-17, oxygen-18, oxygen-12, oxygen-24, and oxygen-15

ELECTRONS – ENERGY LEVELS

• Energy is the capacity to cause a change; the capacity to do work (exert a force on an object).

• The energy level of an electron is the potential energy of that electron.

• Average distance of an electron from the nucleus

• Represented by electron shells

• Shells further from the nucleus represent higher levels of energy

• Energy input (light, etc.) required to move an electron to a higher energy level (further shell)

• Energy released (usually heat) when electron moves back to original shell (ground state)

The higher up the ball is, the more potential energy is has

ELECTRONS – CONFIGURATION/CHEMICAL PROPERTIES• Chemical properties of an element depend on the number of

valence electrons (outermost shell).

• Elements with the same number of valence electrons exhibit similar chemical properties

ELECTRONS – ORBITALS

• Electron Orbital – The area within an atom that electrons have the highest probability of being found.

• Each shell has a specific number of orbitals.

• Each orbital can contain up to 2 electrons.

• Different orbitals represented by different shapes

• Types of orbitals – • s orbitals (spherical) – 1 per shell• p orbitals (dumbbell-shaped) – 3 per shell• d orbitals – 5 per shell

GROUND STATE ELECTRON CONFIGURATIONThe ground state electron configuration for oxygen is…

1s22s22p4

For aluminum…

1s22s22p63s23p1

What is the ground state electron configuration for…

• Sodium (Na)

• Sulfur (S)

• Potassium (K)

• Scandium (Sc)

CHEMICAL BONDING

• Most stable atoms are those with completely filled valence shells (usually eight valence electrons).

• Atoms interact with each other to form bonds in order to fill their valence shells.

• Types of chemical bonding•Covalent• Ionic•Metallic

COVALENT BONDS

Atoms fill their valence shells by sharing electrons with other atoms

Occur between nonmetals

Generally have a low to intermediate difference in electronegativity

Polar covalent bonds – electrons are shared unequally; one atom has a more negative charge (δ-) and the other has a more positive charge (δ+).

Nonpolar covalent bonds – electrons are shared equally and result in a net charge of zero. No difference in electronegativity.

Bond order – ½(# of bonding electrons); single, double, or triple bonds.

Bonding capacity = atom’s valence (number of unpaired valence electrons).

The structure of a molecule depends on the type and number of covalent bonds present, as well as the number of paired electrons.

Tetrahedral

Bent

IONIC BONDS

Atoms that have a high difference in electronegativity

Atom with higher electronegativity strips electron away from atom with lower electronegativity to create two ions.

Cation – positively charged ion

Anion – negatively charged ion

Resulting ions are attracted to each other, forming an ionic bond.

Metals form cations; nonmetals form anions

IONIC COMPOUNDS

Also referred to as salts

Most commonly found in nature as crystals

Formula only represents ratio of atoms (NaCl); individual molecules do not form

Ionic compounds form lattice structures

Cl–

Na+

WEAK CHEMICAL BONDS

Hydrogen bonds – Hydrogen atoms in a polar covalent bond are attracted to another electronegative atom

Van der Waals Interactions – Due to delocalized electrons, molecules form temporary “hotspots” of charge that enable molecules to stick together.

MOLECULAR SHAPE AND FUNCTION

Molecular shape depends on the positions of its atoms’ valence orbitals.

In covalent bonds, overlapping of orbitals results in orbital hybridization.

Specific shapes are associated with specific hybrid orbitals.

MOLECULAR SHAPE AND FUNCTION

Molecular function depends on molecular shape

Molecules interact with each other depending on shape

Enzymes – Fit into specific receptor cites based on shape

CHEMICAL REACTIONS

Molecular bonds break Atoms rearrange themselves New molecular bonds form

CHEMICAL REACTIONS

Chemical Equilibrium – The rate of a chemical reaction is the same as the rate of the reverse reaction

Concentration of reactants and products is NOT necessarily equal when chemical equilibrium is reached.

Reactions still occur after equilibrium is reached, but there is no net change in the concentration of either reactants or products.

EXAM PRACTICE

The atomic number of sulfur is 16. Sulfur combines with hydrogen by covalent bonding to form a compound, hydrogen sulfide. Based on the electron configuration of sulfur, we can predict that the molecular formula of the compound will be;

A.) HS b.) HS2 c.) H2S d.)H3S2

E.) H4S

EXAM PRACTICE

Draw the Lewis Structure for the following molecules…

HCN

H2S

CH2Br2

EXAM PRACTICE

Why is it important that we understand chemistry when we are studying biology?

What coefficients must be placed in the blanks so that all atoms are accounted for in the products?

___C6H12O6 → ___C2H6O + ___CO2

___S8 +  ___O2 →  ___SO2

___Al2(SO4)3 +  ___Ca(OH)2 →  ___Al(OH)3 +  ___CaSO4