2 Chemistry Comes Alive Part A. Matter The “stuff” of the universe Anything that has mass and takes up space States of matter Solid – has definite.

2Chemistry Comes Alive

Part A

Matter

The “stuff” of the universe

Anything that has mass and takes up space

States of matter

Solid – has definite shape and volume

Liquid – has definite volume, changeable shape

Gas – has changeable shape and volume

Energy

The capacity to do work (put matter into motion)

Types of energy

Kinetic – energy in action

Potential – energy of position; stored (inactive) energy

Forms of Energy

Chemical – stored in the bonds of chemical substances

Electrical – results from the movement of charged particles

Mechanical – directly involved in moving matter

Radiant or electromagnetic – energy traveling in waves (i.e., visible light, ultraviolet light, and X rays)

Energy Form Conversions

Energy is easily converted from one form to another

During conversion, some energy is “lost” as heat

Composition of Matter

Elements – unique substances that cannot be broken down by ordinary chemical means

Atoms – more-or-less identical building blocks for each element

Atomic symbol – one- or two-letter chemical shorthand for each element

Properties of Elements

Each element has unique physical and chemical properties

Physical properties – those detected with our senses

Chemical properties – pertain to the way atoms interact with one another

Major Elements of the Human Body

Oxygen (O)

Carbon (C)

Hydrogen (H)

Nitrogen (N)

96% of body matter

Lesser and Trace Elements of the Human Body

Lesser elements make up 3.9% of the body and include:

Calcium (Ca), phosphorus (P), potassium (K), sulfur (S), sodium (Na), chlorine (Cl), magnesium (Mg), iodine (I), and iron (Fe)

Trace elements make up less than 0.01% of the body

They are required in minute amounts, and are found as part of enzymes

Atomic Structure

The nucleus consists of neutrons and protons

Neutrons – have no charge and a mass of one atomic mass unit (amu)

Protons – have a positive charge and a mass of 1 amu

Electrons are found orbiting the nucleus

Electrons – have a negative charge and 1/2000 the mass of a proton (0 amu)

Models of the Atom

Planetary Model – electrons move around the nucleus in fixed, circular orbits

Orbital Model – regions around the nucleus in which electrons are most likely to be found

Models of the Atom

Figure 2.1

Identification of Elements

Atomic number – equal to the number of protons

Mass number – equal to the mass of the protons and neutrons

Atomic weight – average of the mass numbers of all isotopes

Isotope – atoms with same number of protons but a different number of neutrons

Radioisotopes – atoms that undergo spontaneous decay called radioactivity

Identification of Elements

Figure 2.2

Identification of Elements

Figure 2.3

Molecules and Compounds

Molecule – two or more atoms held together by chemical bonds

Compound – two or more different kinds of atoms chemically bonded together

Mixtures and Solutions

Mixtures – two or more components physically intermixed (not chemically bonded)

Solutions – homogeneous mixtures of components

Solvent – substance present in greatest amount

Solute – substance(s) present in smaller amounts

Concentration of Solutions

Percent, or parts per 100 parts

Molarity, or moles per liter (M)

A mole of an element or compound is equal to its atomic or molecular weight (sum of atomic weights) in grams

Colloids and Suspensions

Colloids, or emulsions, are heterogeneous mixtures whose solutes do not settle out

Example: Jello and Cytosol

Suspensions are heterogeneous mixtures with visible solutes that tend to settle out

Example: Blood

Mixtures Compared with Compounds

No chemical bonding takes place in mixtures

Most mixtures can be separated by physical means

Mixtures can be heterogeneous or homogeneous

Compounds cannot be separated by physical means

All compounds are homogeneous

Chemical Bonds

Electron shells, or energy levels, surround the nucleus of an atom

Bonds are formed using the electrons in the outermost energy level

Valence shell – outermost energy level containing chemically active electrons

Octet rule – except for the first shell which is full with two electrons, atoms interact in a manner to have eight electrons in their valence shell

Chemically Inert Elements

Inert elements have their outermost energy level fully occupied by electrons

Figure 2.4a

Chemically Reactive Elements

Reactive elements do not have their outermost energy level fully occupied by electrons

Figure 2.4b

Types of Chemical Bonds

Ionic

Covalent

Hydrogen

Ionic Bonds

Ions are charged atoms resulting from the gain or loss of electrons

Anions have gained one or more electrons

Cations have lost one or more electrons

Formation of an Ionic Bond

Ionic bonds form between atoms by the transfer of one or more electrons

Ionic compounds form crystals instead of individual molecules

Example: NaCl (sodium chloride)

Formation of an Ionic Bond

Figure 2.5a

Formation of an Ionic Bond

Figure 2.5b

Covalent Bonds

Covalent bonds are formed by the sharing of two or more electrons

Electron sharing produces molecules

Single Covalent Bonds

Figure 2.6a

Double Covalent Bonds

Figure 2.6b

Triple Covalent Bonds

Figure 2.6c

Polar and Nonpolar Molecules

Electrons shared equally between atoms produce nonpolar molecules

Unequal sharing of electrons produces polar molecules

Atoms with six or seven valence shell electrons are electronegative

Atoms with one or two valence shell electrons are electropositive

Figure 2.8

Comparison of Ionic, Polar Covalent, and Nonpolar Covalent Bonds

Hydrogen Bonds

Too weak to bind atoms together

Common in dipoles such as water

Responsible for surface tension in water

Important as intramolecular bonds, giving the molecule a three-dimensional shape

Hydrogen Bonds

Figure 2.9

Chemical Reactions

Occur when chemical bonds are formed, rearranged, or broken

Are written in symbolic form using chemical equations

Chemical equations contain:

Number and type of reacting substances, and products produced

Relative amounts of reactants and products

Examples of Chemical Reactions

Patterns of Chemical Reactions

Combination reactions: Synthesis reactions which always involve bond formation

A + B AB

Decomposition reactions: Molecules are broken down into smaller molecules

AB A + B

Exchange reactions: Bonds are both made and broken

AB + C AC + B

Oxidation-Reduction (Redox) Reactions

Reactants losing electrons are electron donors and are oxidized

Reactants taking up electrons are electron acceptors and become reduced

Therefore, both decomposition and electron exchange occur.

Energy Flow in Chemical Reactions

Exergonic reactions – reactions that release energy

Usually when a bond is broken.

Endergonic reactions – reactions whose products contain more potential energy than did its reactants

Reversibility in Chemical Reactions

All chemical reactions are theoretically reversible

A + B AB

AB A + B

If neither a forward nor reverse reaction is dominant, chemical equilibrium is reached

Factors Influencing Rate of Chemical Reactions

Temperature – chemical reactions proceed quicker at higher temperatures

Particle size – the smaller the particle the faster the chemical reaction

Concentration – higher reacting particle concentrations produce faster reactions

Catalysts – increase the rate of a reaction without being chemically changed

Enzymes – biological catalysts

2Chemistry Comes Alive:

BiochemistryPart B

Biochemistry

Inorganic compounds

Do not contain carbon

Water, salts, and many acids and bases

Organic compounds

Contain carbon, are covalently bonded, and are often large

Inorganic: Water

High heat capacity – absorbs and releases large amounts of heat before changing temperature

High heat of vaporization – changing from a liquid to a gas requires large amounts of heat

Polar solvent properties – dissolves ionic substances, forms hydration layers around large charged molecules, and serves as the body’s major transport medium

Inorganic: Water

Reactivity – is an important part of hydrolysis and dehydration synthesis reactions

Cushioning – resilient cushion around certain body organs

Inorganic: Salts

Inorganic compounds

Contain cations other than H+ and anions other than OH–

Are electrolytes; they conduct electrical currents

Inorganic: Acids and Bases

Acids release H+ and are therefore proton donors

HCl H+ + Cl –

Bases release OH– and are proton acceptors

NaOH Na+ + OH–

Inorganic: Acid-Base Concentration (pH)

Acidic solutions have higher H+ concentration and therefore a lower pH

Alkaline solutions have lower H+ concentration and therefore a higher pH

Neutral solutions have equal H+ and OH– concentrations

Inorganic: Acid-Base Concentration (pH)

Acidic: pH 0–6.99

Basic: pH 7.01–14

Neutral: pH 7.00

Figure 2.12

Inorganic: Buffers

Systems that resist abrupt and large swings in the pH of body fluids

Carbonic acid-bicarbonate system

Carbonic acid dissociates, reversibly releasing bicarbonate ions and protons

The chemical equilibrium between carbonic acid and bicarbonate resists pH changes in the blood

Organic Compounds

Molecules unique to living systems contain carbon and hence are organic compounds

They include:

Carbohydrates

Lipids

Proteins

Nucleic Acids

Organic: Carbohydrates

Figure 2.13a

Contain carbon, hydrogen, and oxygen

Their major function is to supply a source of cellular food

Examples:

Monosaccharides or simple sugars

Organic: Carbohydrates

Figure 2.13b

Disaccharides or double sugars

Organic: Carbohydrates

Figure 2.13c

Polysaccharides or polymers of simple sugars

Organic: Lipids

Contain C, H, and O, but the proportion of oxygen in lipids is less than in carbohydrates

Examples:

Neutral fats or triglycerides

Phospholipids

Steroids

Eicosanoids

Organic: Neutral Fats (Triglycerides)

Figure 2.14a

Composed of three fatty acids bonded to a glycerol molecule

Organic: Other Lipids

Figure 2.14b

Phospholipids – modified triglycerides with two fatty acid groups and a phosphorus group

Organic: Other Lipids

Figure 2.14c

Steroids – flat molecules with four interlocking hydrocarbon rings

Eicosanoids – 20-carbon fatty acids found in cell membranes

Organic: Representative Lipids Found in the Body

Neutral fats – found in subcutaneous tissue and around organs

Phospholipids – chief component of cell membranes

Steroids – cholesterol, bile salts, vitamin D, sex hormones, and adrenal cortical hormones

Fat-soluble vitamins – vitamins A, E, and K

Eicosanoids – prostaglandins, leukotriens, and thromboxanes

Lipoproteins – transport fatty acids and cholesterol in the bloodstream

Organic: Amino Acids

Building blocks of protein, containing an amino group and a carboxyl group

Amino acid structure

Organic: Amino Acids

Figure 2.15a-c

Organic: Amino Acids

Figure 2.15d, e

Organic: Protein

Figure 2.16

Macromolecules composed of combinations of 20 types of amino acids bound together with peptide bonds

Organic: Structural Levels of Proteins

Primary – amino acid sequence

Secondary – alpha helices or beta pleated sheets

Organic: Structural Levels of Proteins

Tertiary – superimposed folding of secondary structures

Quaternary – polypeptide chains linked together in a specific manner

Organic: Structural Levels of Proteins

Figure 2.17a-c

Organic: Structural Levels of Proteins

Figure 2.17d, e

Organic: Fibrous and Globular Proteins

Fibrous proteins

Extended and strandlike proteins

Examples: keratin, elastin, collagen, and certain contractile fibers

Globular proteins

Compact, spherical proteins with tertiary and quaternary structures

Examples: antibodies, hormones, and enzymes

Organic: Protein Denuaturation

Figure 2.18a

Reversible unfolding of proteins due to drops in pH and/or increased temperature

Organic: Protein Denuaturation

Figure 2.18b

Irreversibly denatured proteins cannot refold and are formed by extreme pH or temperature changes

Organic: Molecular Chaperones (Chaperonins)

Help other proteins to achieve their functional three-dimensional shape

Maintain folding integrity

Assist in translocation of proteins across membranes

Promote the breakdown of damaged or denatured proteins

Organic: Characteristics of Enzymes

Most are globular proteins that act as biological catalysts

Holoenzymes consist of an apoenzyme (protein) and a cofactor (usually an ion)

Enzymes are chemically specific

Frequently named for the type of reaction they catalyze

Enzyme names usually end in -ase

Lower activation energy

Organic: Characteristics of Enzymes

Figure 2.19

Organic: Mechanism of Enzyme Action

Enzyme binds with substrate

Product is formed at a lower activation energy

Product is released

Enzyme-substrate

complex (E–S)

1

2

3

Internal rearrangements leading to catalysis

Free enzyme (E)

Active site

Enzyme (E) Substrates (s)

Amino acids

H20

Peptide bond

Dipeptide product (P)

Organic: Mechanism of Enzyme Action

Figure 2.20

Organic: Nucleic Acids

Composed of carbon, oxygen, hydrogen, nitrogen, and phosphorus

Their structural unit, the nucleotide, is composed of N-containing base, a pentose sugar, and a phosphate group

Five nitrogen bases contribute to nucleotide structure – adenine (A), guanine (G), cytosine (C), thymine (T), and uracil (U)

Two major classes – DNA and RNA

Organic: Deoxyribonucleic Acid (DNA)

Double-stranded helical molecule found in the nucleus of the cell

Replicates itself before the cell divides, ensuring genetic continuity

Provides instructions for protein synthesis

Organic: Structure of DNA

Figure 2.21a

Organic: Structure of DNA

Figure 2.21b

Organic: Ribonucleic Acid (RNA)

Single-stranded molecule found in both the nucleus and the cytoplasm of a cell

Uses the nitrogenous base uracil instead of thymine

Three varieties of RNA: messenger RNA, transfer RNA, and ribosomal RNA

Organic: Adenosine Triphosphate (ATP)

Source of immediately usable energy for the cell

Adenine-containing RNA nucleotide with three phosphate groups

Organic: Adenosine Triphosphate (ATP)

Figure 2.22

Organic: How ATP Drives Cellular Work

Figure 2.23