Chapter 2 - Chemistry come alive

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• Matter is anything that has mass and occupies space.

States of Matter

• Solid possesses a definite shape and volume.
• Liquid has a definite volume but a changeable shape.
• Gas exhibits a changeable shape and volume.

Energy

• Energy is the capacity to perform work, which involves putting matter into motion.

Types of Energy

• Kinetic energy pertains to energy in action.
• Potential energy is the energy of position, stored in an inactive state.
• Energy can be transferred from potential to kinetic energy.

Forms of Energy

• Chemical energy is stored within the bonds of chemical substances.
• Electrical energy arises from the movement of charged particles.
• Mechanical energy is directly involved in the movement of matter.
• Radiant or electromagnetic energy travels in waves, like visible light, ultraviolet light, and X-rays.

Energy Form Conversions

• Energy can be easily converted from one form to another.
• Energy conversion is inefficient.
• Some energy is "lost" as heat during conversion processes and becomes partly unusable.

Composition of Matter

• Elements are unique substances that cannot be broken down by ordinary chemical means.
• Matter is composed of elements.
• Atoms serve as the nearly identical building blocks for each element.
• Atomic symbols are one- or two-letter chemical shorthands representing each element.

Properties of Elements

• Each element has specific physical and chemical traits.
• Physical properties are those detectable through our senses.
• Chemical properties relate to how atoms interact (bond).

Major Elements of the Human Body

• Four elements constitute 96.1% of body mass.
• Oxygen's symbol is O.
• Carbon's symbol is C.
• Hydrogen's symbol is H.
• Nitrogen's symbol is N.

Lesser and Trace Elements of the Human Body

• Lesser elements constitute 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 comprise less than 0.01% of the body, with 11 elements making up the group
• Trace elements are needed in small quantities and are components of enzymes, which they activate.
• Chromium (Cr) and copper (Cu) are examples of trace elements.

Atomic Structure

• Atoms are composed of subatomic particles: protons, neutrons, and electrons.
• The nucleus of an atom contains neutrons and protons.
• Neutrons have no charge and possess a mass of one atomic mass unit (amu).
• Protons carry a positive charge and have a mass of 1 amu.
• Electrons orbit the nucleus within an electron cloud.
• Electrons are negatively charged and have a mass of approximately 1/2000 of a proton (0 amu).
• The number of protons and electrons is always equal in an atom
• The nucleus contains almost the entire mass of the atom

Identification of Elements

• Each element contains a different number of subatomic particles
• Atomic number corresponds to the number of protons.
• Mass number is equal to the mass of protons and neutrons in the nucleus.
• Written as superscript to the left of the atomic symbol
• Atomic weight is the average of the mass numbers of all isotopes of an element.
• Isotopes are structural variations of atoms, that have the same number of protons but differ in their number of neutrons.
• Radioisotopes are heavy isotopes that decompose to more stable forms, and this spontaneous decay is called radioactivity. Radioisotopes can transform to a different element and can be detected with scanners.

Molecules and Compounds

• Most atoms chemically combine with other atoms to form molecules and compounds
• Molecules consist of two or more atoms held together by chemical bonds: H2 or C6H12O6.
• Compounds are substances made up of two or more different types of atoms chemically bonded: C6H12O6 but not H2

Mixtures and Solutions

• Most matter exists as mixtures
• Mixtures involve two or more components that are physically intermixed, not chemically bonded.
• There are three types of mixtures; solutions, colloids, and suspensions.
• Solutions are homogeneous mixtures of components that are made of gases, liquids, or solids dissolved in water, most are actual solutions in the body.
• The solvent is the substance present in the greatest amount within a solution, often a liquid.
• The solute is present in smaller amounts.

Chemical Bonds

• Electron shells, or energy levels, surround the nucleus of an atom.
• The electrons can occupy up to seven electron shells around the nucleus
• Bonds form through the interactions of electrons in the outermost energy level.
• The valence shell refers to the outermost energy level, containing chemically active electrons.
• Electrons in the valence shell have the most potential energy
• The octet rule dictates that atoms interact to have eight electrons in their valence shell, except for the innermost shell, which is stable with two electrons.

Chemically Inert Elements

• Inert elements fully occupy their outermost energy level with electrons, making them non-reactive, they are stable and comprise the noble gasses.

Chemically Reactive Elements

• Reactive elements do not have their outermost energy level completely filled with electrons, causing them to react with other elements.
• Chemically Reactive Elements will tend to gain, loss, or share electrons with other atoms to achieve stability

Types of Chemical Bonds

• Ionic bonds
• Covalent bonds
• Hydrogen bonds

Ionic Bonds

• Ions are charged atoms that result from the gain or loss of electrons.
• Anions gain one or more electrons and have a negative charge.
• Cations lose one or more electrons and have a positive charge.

Formation of an Ionic Bond

• Ionic bonds form through the transfer of one or more electrons between atoms.
• Salt (NaCl), or sodium chloride, provides a common example of ionic bonding.

Covalent Bonds

• Covalent bonds are created when two or more atoms share electrons.
• Allows each atom to fill its valence shell and fill at least part of the time
• Electron sharing creates molecules.

Polar and Nonpolar Molecules

• Electrons are equally shared between atoms in Nonpolar molecules, creating nonpolar molecules and resulting in equal charge distribution.
• Unequal sharing of electrons produces polar molecules, creating regions of partial positive and negative charge.
• Atoms containing six or seven valence shell electrons are electronegative (e.g., oxygen) and tend to attract electrons more strongly in a chemical bond.
• Atoms containing one or two valence shell electrons are electropositive (e.g., sodium).

Hydrogen Bonds

• Hydrogen bonds are attractive forces that are too weak to bind atoms together permanently, they are not true bonds.
• Hydrogen bonds are commonly found in dipoles such as water molecules.
• They are responsible for surface tension in water.
• Act as intramolecular bonds, holding a large molecule in a three-dimensional shape.

Properties of Water

• High heat capacity allows water to absorb and release significant amounts of heat before changing temperature.
• High heat of vaporization indicates it requires a lot of heat to change from liquid to a gas.
• Polar solvent properties enable it to dissolve ionic substances and to form hydration layers around large charged molecules, also allowing water to function as the body's major transport medium.
• Water's reactivity is essential for hydrolysis and dehydration synthesis reactions.
• Water's cushioning provides protection around some organs in the body.

Chemical Reactions

• Chemical reactions occur when chemical bonds form, rearrange, or break.
• These reactions are written in symbolic form using chemical equations.
• A subscript indicates atoms joined by bonds, a prefix indicated the number of unjoined atoms or molecules.
• Chemical equations include the number and types of reacting substances and the produced compounds. Reactants and products.
• Chemical equations include the relative quantities of reactants and products.

Patterns of Chemical Reactions

• Combination or synthesis reactions always involve forming chemical bonds to create larger molecules from smaller ones: A + B → AB, also known as Anabolic.
• Decomposition reactions involve breaking down larger compounds into smaller molecules: AB → A + B, also known as Catabolic.
• Exchange reactions involve both making and breaking chemical bonds, with some elements or molecules being exchanged: AB + C → AC + B, also known as displacement reactions.

Energy Flow in Chemical Reactions

• All chemical reactions are either Exergonic or endergonic.
• Exergonic reactions net release energy.
• Products have less potential then reactants.
• These reactions are catabolic and oxidative reactions.
• Endergonic reactions net absorb energy.
• Products have more potential energy the reactants.
• These reaction are anabolic reactions

Factors Influencing Rate of Chemical Reactions

• Increased temperature typically quickens chemical reactions.
• Chemical reactions proceed quicker with decreased size of the particle.
• Higher concentration of reactants generally yields more rapid reactions.
• Catalysts increase reaction rates without undergoing chemical changes.
• Enzymes serve as biological catalysts to facilitate specific biochemical reactions in the body.

Biochemistry

• Organic compounds contain carbon, are covalently bonded, and are often large.
• Inorganic compounds do not contain carbon, examples include water, salts, and many acids and bases.

Acids and Bases

• Acids release H+ and are therefore proton donors; HCl → H+ + Cl-.
• Bases release OH- and are proton acceptors; NaOH → Na+ + OH-.

Acid-Base Concentration (pH)

• Relative free (H+) of a solution measured on pH scale
• As free [H+] increases, acidity increases
• [OH-] decreases as H+ increases.
• pH decreases As free [H+] decreases alkalinity increases [OH-] increases as H+ increases.
• PH increases
• Acidic solutions have higher H+ concentration and therefore a lower pH level, pH being between 0–6.99.
• Alkaline solutions have lower H+ concentration and therefore a higher pH level, pH being between 7.01–14.
• Neutral solutions have equal H+ and OH- concentrations and have the level of pH 7.00.

Buffers

• Buffers are systems that resist abrupt and large swings in the pH of body fluids.
• Acidity reflects only free H+ in solution.
• The Carbonic acid-bicarbonate system sees carbonic acid dissociate, reversibly releasing bicarbonate ions and protons.
• 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, and include carbohydrates, lipids, proteins, and nucleic acids.
• Carbon is electroneutral
• Shares electrons never gains or loses them.
• Forms four covalent bonds with other elements.

Carbohydrates

• Contain carbon, hydrogen, and oxygen.
• Their major function is to supply a source of cellular food.
• Examples are Monosaccharides (simple sugars), Disaccharides (double sugars), and Polysaccharides (polymers of simple sugars).

Lipids

• Contain carbon, hydrogen, and oxygen but the proportion of oxygen in lipids is less than in carbohydrates.
• Lipids may sometimes contain phosphorus
• Insoluble in water
• Examples are Neutral fats or triglycerides, phospholipids, and steroids.

Neutral Fats (Triglycerides)

• Composed of three fatty acids bonded to a glycerol molecule.

Other Lipids

• Phospholipids are modified triglycerides with two fatty acid groups and a phosphorus group.
• Steroids are flat molecules with four interlocking hydrocarbon rings.
• Cholerole most important steroid.

Representative Lipids Found in the Body

• Neutral fats are found in subcutaneous tissue and around organs.
• Phospholipids are the chief component of cell membranes.
• Steroids include cholesterol, sex hormones, and adrenal cortical hormones.
• Fat-soluble vitamins are vitamins A, D, E, and K.
• Lipoproteins transport fatty acids and cholesterol in the bloodstream.

Amino Acids

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

Protein

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

Structural Levels of Proteins

• Primary: amino acid sequence.
• Secondary: alpha helices or beta-pleated sheets.
• Tertiary: superimposed folding of secondary structures.
• Quaternary: polypeptide chains linked together in a specific manner.

Fibrous and Globular Proteins

• Fibrous proteins are extended and strandlike proteins, examples include keratin, elastin, collagen, and certain contractile fibers.
• Globular proteins are compact, spherical proteins with tertiary and quaternary structures, examples include antibodies, hormones, and enzymes.
• Active sites destroyed.

Protein Denaturation

• Reversible unfolding of proteins due to drops in pH and/or increased temperature.
• Irreversibly denatured proteins cannot refold and are formed by extreme pH or temperature changes.

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 and act on specific substrate
• Frequently named for the type of reaction they catalyze, and enzyme names usually end in -ase.
• Enzymes lower activation energy.

Mechanism of Enzyme Action

• Enzyme binds with the substrate.
• Product is formed at a lower activation energy.
• Product is released.

Nucleic Acids

• Composed of carbon, oxygen, hydrogen, nitrogen, and phosphorus.
• Their structural unit, the nucleotide, is composed of an 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.

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.
• Double-stranded helical molecule (double helix) in the cell nucleus.

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.

Adenosine Triphosphate (ATP)

• Source of immediately usable energy for the cell.
• Adenine-containing RNA nucleotide with three phosphate groups.The chemical energy in glucose is captured in this important molecule.
• Directly powers chemical reactions in cells.
• Energy from immediately useable by all body cells.