Chapter 2

Atoms and Molecules: The Chemical Basis of Life

Inorganic vs. Organic

Inorganic compounds - simple substances that do not contain carbon Ex. water

Organic compounds – carbon-containing compounds that are large and complex Ex. Glucose

Organic compounds all contain:

6%0%

94%

0%

1. Oxygen2. Nitrogen3. Carbon4. Zinc

Elements

Elements – substances that cannot be broken down into simpler substances

4 elements responsible for more than 96% of the mass of most organisms: Oxygen Carbon Hydrogen Nitrogen

See table 2-1 in book p.24

Other elements that make up living organisms

Calcium Phosphorus Potassium Magnesium Sodium Iron Sulfur

What is the most abundant element in the human body?

29%

12%

53%

6%

1. Oxygen 2. Carbon3. Hydrogen4. Nitrogen

Chemistry Quick Review

Atom – the smallest portion of an element that retains its chemical properties

Subatomic particles: Electron – negative charge Proton – positive charge Neutron – uncharged

# of electrons = # of protons Nucleus – protons and neutrons Electrons – move rapidly through space

around nucleus

Atomic Number

Each kind of element has a fixed number of protons in the atomic nucleus

Written as a subscript to the left of the chemical symbol Example: 8O Oxygen nucleus contains 8 protons Determines the atom’s identity and

defines the element

Which element has an atomic number of 7?

0% 0%

100%

0%

1. Hydrogen2. Carbon3. Nitrogen4. Helium

The Periodic Table

Chart in which elements are arranged in order by atomic number

Can be used to determine electron configurations Bohr model – shows the electrons

arranged in a series of concentric circles around the nucleus

Bohr Model

What element is this?

Atomic Mass Mass of protons + neutrons Mass of electron = 1/1800 the mass of

a proton or neutron Atomic mass number is a superscript

to the left of the symbol Example: 16O

Isotopes

Atom with different number of neutrons (different masses)

Most elements mixture of isotopes Ex. Carbon-12, Carbon-14

Mass of element is average of the masses of its isotopes Atomic mass of Carbon = 12.011

Isotopes

Radioisotopes – unstable isotopes Tend to break down (decay) to a more

stable isotope Emit radiation when they decay Ex. Carbon-14 decays to Nitrogen

Electrons move in orbitals

Orbitals are more like “electron clouds”

The farther away from the nucleus, the more energy the electrons have

Valence electrons – the most energetic electrons Occupy valence (outer) shell

Chemical Reactions

Valence electrons participate in chemical reactions

When valence shell is full, it is stable

When valence shell is not full, atoms tend to lose, gain, or share electrons

To be full, the first electron shell has how many electrons?

0%

94%

0%6%

0%

1. 12. 23. 44. 85. 18

Compounds and Molecules

Atoms combine to form compounds and molecules

Compounds - 2 or more different elements combined in a fixed ratio Ex. NaCl (table salt)

Molecules - 2 or more atoms combine chemically Ex. O2, DNA

Are all molecules compounds?

71%

29%

1. Yes2. No

Molecule or Compound? O2

65%

24%

12%

1. Molecule2. Compound3. Both

Chemical Formulas

Represents chemical composition Simplest formula – most simple ratio

Ex. NH2

Molecular formula – actual numbers of each type of atom per molecule Ex. N2H4

Structural formula – shows arrangement of atoms Ex. Water H – O – H

Chemical Equations

Reactants – participate in reaction Products – formed by the reaction Example – cellular respiration C6H12O6 + 6O2 -> 6CO2 + 6H2O + Energy

Chemical Bonds

Valence electrons dictate # of bonds

2 types of chemical bonds: Covalent – atoms share electrons Ionic - attraction between positive

cations and negative anions Transfer electrons

Covalent Bonds

Ex. H2 gas Each atom has 1 electron 2 electrons fill valence shell Both atoms attract the electrons

(share) Valence shell is full w/ 2 electrons

Types of Covalent Bonds

Single covalent bond – 1 pair of electrons is shared

Double covalent bond – 2 pairs of electrons shared

Triple covalent bond – 3 pairs of electrons shared

Covalent Bonds

Electronegativity - measure of atom’s attraction for shared electrons in chemical bonds

Oxygen, Nitrogen, Fluorine, Chlorine very electronegative

Can be polar or nonpolar Similar electronegativities = nonpolar bonds Different electronegativities = polar bonds

Electrons are pulled closer to the nucleus of the atom with the higher electronegativity

Polar Molecules

Molecule with one or more polar covalent bonds

One end with a partial positive charge and other end a partial negative charge

Ex. Water (p.31)

A water molecule is polar because

24%

0%

76%1. The electrons orbit

the H atoms more closely

2. The electrons orbit the O atom more closely

3. The electrons orbit all atoms equally

Ionic Bonds

Ionic compound – consists of anions and cations bonded together

Ex. NaCl (p.31 & 32) Na – 1 valence electron Cl – 7 valence electrons Cl takes electron from Na to complete

valence shell

Hydrogen Bonds

Weak attractions Important in determining the 3-D

structure of large molecules DNA Proteins

Why are hydrogen bonds essential to the function of DNA?

33% 33%33%1. They keep the 2

strands tightly bonded together

2. They allow the 2 strands to separate for replication

3. They are strong bonds

Redox Reactions

Reaction that involves electron transfer

Cellular Respiration and Photosynthesis

Oxidation – atom/ion loses electron Reduction – atom/ion gains electron

Water

70% of total body weight Reactant/product in many chemical

reactions Solvent for most biological reactions

Hydrophilic – react with water Hydrophobic – not disrupted/dissolved

by water

Which of the following substances is hydrophobic?

Salt

Suga

r O

il

33% 33%33%1. Salt2. Sugar3. Oil

Properties of Water

Cohesive – water molecules stick to each other

Adhesive – water molecules stick to other substances

Capillary action – cohesion and adhesion working together Water will move against gravity in a

narrow tube In plants, water moves from soil to roots

Properties of Water

Surface tension – water molecules crowd together at the surface strong layer

High specific heat Maintains a stable temperature

High heat of vaporization Much heat required to change to water

vapor

Acids and Bases

Acids – proton donors Acid -> H+ + Anion Acidic solutions have higher hydrogen

ion concentration Turn blue litmus paper red Sour taste HCl – inorganic acid Acetic Acid – from vinegar, Lactic Acid –

from sour milk (organic acids)

Acids and Bases

Bases – proton acceptors Base -> OH- + Cation Basic solutions have lower hydrogen

ion concentration Turn red litmus paper blue Feel slippery to the touch Ex. Sodium Hydroxide, Ammonia –

inorganic Purine and Pyrimidine – organic

pH scale

Logarithmic expression of the hydrogen ion concentration of solution

7 = neutral Below 7 = Acid Above 7 = Base