© 2016 Pearson Education, Inc.

Chemistry and Physiological Reactions
Body is made up of many chemicals Chemistry underlies all physiological reactions: Movement, digestion, pumping of heart, nervous system Chemistry can be broken down into: Basic chemistry Biochemistry © 2016 Pearson Education, Inc.

Part 1 – Basic Chemistry 2.1 Matter and Energy Matter
Matter is anything that has mass and occupies space Matter can be seen, smelled, and/or felt Weight is mass plus the effects of gravity © 2016 Pearson Education, Inc.

Matter States of matter Matter can exist in three possible states:
Solid: definite shape and volume Liquid: changeable shape; definite volume Gas: changeable shape and volume © 2016 Pearson Education, Inc.

Energy Energy is the capacity to do work or put matter into motion
Energy does not have mass, nor does it take up space The greater the work done, the more energy it uses up © 2016 Pearson Education, Inc.

Energy (cont.) Kinetic versus potential energy
Energy exists in two possible forms: Kinetic – energy in action Potential – stored (inactive) energy Energy can be transformed from potential to kinetic energy Stored energy can be released, resulting in action © 2016 Pearson Education, Inc.

Energy (cont.) Forms of energy Chemical energy Electrical energy
Stored in bonds of chemical substances Electrical energy Results from movement of charged particles Mechanical energy Directly involved in moving matter Radiant or electromagnetic energy Travels in waves (example: heat, visible light, ultraviolet light, and X rays) © 2016 Pearson Education, Inc.

Energy (cont.) Energy form conversions
Energy may be converted from one form to another Example: turning on a lamp converts electrical energy to light energy Energy conversion is inefficient Some energy is “lost” as heat, which can be partly unusable energy © 2016 Pearson Education, Inc.

2.2 Atoms and Elements All matter is composed of elements
Elements are substances that cannot be broken down into simpler substances by ordinary chemical methods Four elements make up 96% of body: Carbon, oxygen, hydrogen, and nitrogen 9 elements make up 3.9% of body 11 elements make up <0.01% Periodic table lists all known elements © 2016 Pearson Education, Inc.

2.2 Atoms and Elements All elements are made up of atoms, which are:
Unique building blocks for each element Smallest particles of an element with properties of that element What give each element its particular physical & chemical properties © 2016 Pearson Education, Inc.

2.2 Atoms and Elements Atomic symbol
One- or two-letter chemical shorthand for each element Example: “O” for oxygen, “C” for carbon Some symbols come from Latin names: “Na” (natrium) is sodium; “K” (kalium) is potassium © 2016 Pearson Education, Inc.

Table 2.1-1 Common Elements Composing the Human Body
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Table 2.1-2 Common Elements Composing the Human Body (continued)

Table 2.1-3 Common Elements Composing the Human Body (continued)

Structure of Atoms Atoms are composed of three subatomic particles:
Protons Carry a positive charge (+) Weigh an arbitrary 1 atomic mass unit (1 amu) Neutrons Have no electrical charge (0) Also weigh 1 amu Electrons Carry a negative charge () Are so tiny they have virtually no weight (0 amu) © 2016 Pearson Education, Inc.

Structure of Atoms (cont.)
Number of positive protons is balanced by number of negative electrons, so atoms are electrically neutral Protons and neutrons are found in a centrally located nucleus; electrons orbit around the nucleus Chemists devise models of how subatomic particles are put together Planetary model Orbital model © 2016 Pearson Education, Inc.

Structure of Atoms (cont.)
Planetary model: simplified and outdated because it incorrectly depicts electrons in orbits, fixed circular paths Still useful for illustrations Orbital model: current model used that depicts orbitals, probable regions where an electron is most likely to be located (rather than fixed orbits) Shading in regions of greatest electron density results in an electron cloud around nucleus Useful for predicting chemical behavior of atoms © 2016 Pearson Education, Inc.

Figure 2.1 Two models of the structure of an atom.
Nucleus Nucleus Helium atom Helium atom 2 protons (p+) 2 protons (p+) 2 neutrons (n0) 2 neutrons (n0) 2 electrons (e−) 2 electrons (e−) Planetary model Orbital model Proton Neutron Electron Electron cloud © 2016 Pearson Education, Inc.

Identifying Elements Different elements contain different numbers of subatomic particles Hydrogen has 1 proton, 0 neutrons, and 1 electron Helium has 2 protons, 2 neutrons, and 2 electrons Lithium has 3 protons, 4 neutrons, and 3 electrons Identifying facts about an element include its atomic number, mass number, isotopes, and atomic weight © 2016 Pearson Education, Inc.

Figure 2.2 Atomic structure of the three smallest atoms.
Proton Neutron Electron Hydrogen (H) (1p+; 0n0; 1e−) Helium (He) (2p+; 2n0; 2e−) Lithium (Li) (3p+; 4n0; 3e−) © 2016 Pearson Education, Inc.

Identifying Elements (cont.)
Atomic number Number of protons in nucleus Written as subscript to left of atomic symbol Example: 3Li Mass number Total number of protons and neutrons in nucleus Total mass of atom Written as superscript to left of atomic symbol Example: 7Li © 2016 Pearson Education, Inc.

Identifying Elements (cont.)
Isotopes Structural variations of same element Atoms contain same number of protons but differ in the number of neutrons they contain Atomic numbers are same, but mass numbers different Atomic weight Average of mass numbers of all isotope forms of an atom © 2016 Pearson Education, Inc.

Figure 2.3 Isotopes of hydrogen.
Proton Neutron Electron Hydrogen (1H) (1p+; 0n0; 1e−) Deuterium (2H) (1p+; 1n0; 1e−) Tritium (3H) (1p+; 2n0; 1e−) © 2016 Pearson Education, Inc.

Radioisotopes Radioisotopes are isotopes that decompose to more stable forms Atom loses various subatomic particles Sometimes loss results in an isotope becoming a different element As isotope decays, subatomic particles that are being given off release a little energy This energy is referred to as radioactivity Can be detected and measured with scanners © 2016 Pearson Education, Inc.

Radioisotopes (cont.) Radioisotopes are a valuable tool for biological research and medicine Share same chemistry as their stable isotopes so will be taken up by body Can then be used for diagnosis of disease All radioactivity can damage living tissue Some types can be used to destroy localized cancers Some types cause cancer Radon from uranium decay causes lung cancer © 2016 Pearson Education, Inc.

2.3 Combining Matter Molecules and Compounds
Most atoms chemically combine with other atoms to form molecules and compounds Molecule: general term for 2 or more atoms bonded together Compound: specific molecule that has 2 or more different kinds of atoms bonded together Example: C6H12O6 Molecules with only one type of atom (H2 or O2) are just called molecules © 2016 Pearson Education, Inc.

Figure 2.4 The three basic types of mixtures.
Solution Colloid Suspension Solute particles are very tiny, do not settle out or scatter light. Solute particles are larger than in a solution and scatter light; do not settle out. Solute particles are very large, settle out, and may scatter light. Solute particles Solute particles Solute particles Example Example Example Mineral water Jell-O Blood Plasma Settled red blood cells Unsettled Settled © 2016 Pearson Education, Inc.

Mixtures (cont.) Solutions
Are homogeneous mixtures, meaning particles are evenly distributed throughout Solvent: substance present in greatest amount Usually a liquid, such as water Solute(s): substance dissolved in solvent Present in smaller amounts Example: blood sugar – glucose is solute, and blood (plasma) is solvent © 2016 Pearson Education, Inc.

Mixtures (cont.) Solutions (cont.)
True solutions are usually transparent Example: air (gas solution), salt solution, sugar solution Most solutions in body are true solutions of gases, liquids, or solids dissolved in water © 2016 Pearson Education, Inc.

Mixtures (cont.) Concentration of true solutions
Three common ways to express concentrations: Percent of solute in total solution How many parts of solute are in 100 total parts of solution Solvent is usually water Example: 10 parts salt to 90 parts water is a 10% salt solution Milligrams per deciliter (mg/dl) Deciliter equals 1/100th of a liter Example: normal fasting blood glucose levels are around 80 mg/dl © 2016 Pearson Education, Inc.

Mixtures (cont.) Molarity (M) is number of moles of solute per liter of solvent (water) 1 mole of a compound is equal to its molecular weight (sum of atomic weights) in grams Example: glucose (C6H12O6 ) has a molecular wt of amu, so grams of glucose added to enough H2O to make 1 liter is a 1 M solution of glucose 1 mole of any substance always contains 6.02  1023 molecules of that substance This number is called Avogadro’s number Molarities in the body are so small (can be M), they are expressed in millimoles (mM) so mM  1 M © 2016 Pearson Education, Inc.

Mixtures (cont.) Colloids
Also known as emulsions; are heterogeneous mixtures, meaning that particles are not evenly distributed throughout mixture Can see large solute particles in solution, but these do not settle out Gives solution a cloudy or milky look Some undergo sol-gel (solution to gel) transformations Example: Jell-O goes from liquid to gel Cytosol of cell is also a sol-gel type solution © 2016 Pearson Education, Inc.

Mixtures (cont.) Suspensions
Heterogeneous mixtures that contain large, visible solutes that do settle out Example: mixture of water and sand Blood is considered a suspension because if left in a tube, the blood cells will settle out © 2016 Pearson Education, Inc.

Difference Between Mixtures and Compounds
Three main differences: Unlike compounds, mixtures do not involve chemical bonding between components Mixtures can be separated by physical means, such as straining or filtering; compounds can be separated only by breaking their chemical bonds Mixtures can be heterogeneous or homogeneous; compounds are only homogeneous © 2016 Pearson Education, Inc.

2.4 Chemical Bonds Chemical bonds are “energy relationships” between electrons of reacting atoms Chemical bonds are not actual physical structures Electrons are the subatomic particles that are involved in all chemical reactions They determine whether a chemical reaction will take place and if so, what type of chemical bond is formed © 2016 Pearson Education, Inc.

Role of Electrons in Chemical Bonding
Electrons can occupy areas around nucleus called electron shells Each shell contains electrons that have a certain amount of kinetic and potential energy, so shells are also referred to as energy levels Depending on its size, an atom can have up to 7 electron shells Shells can hold only a specific number of electrons; the shell closest to nucleus is filled first Shell 1 can hold only 2 electrons Shell 2 holds a maximum of 8 electrons Shell 3 holds a maximum of 18 electrons © 2016 Pearson Education, Inc.

Role of Electrons in Chemical Bonding (cont.)
Outermost electron shell is called valence shell Electrons in valence shell have the most potential energy because they are farthest from nucleus These are electrons that are involved in chemical reactions © 2016 Pearson Education, Inc.

Role of Electrons in Chemical Bonding (cont.)
Octet rule (rule of eights) Atoms desire 8 electrons in their valence shell Exceptions: smaller atoms (examples: H and He) want only 2 electrons in shell 1 Desire to have 8 electrons is driving force behind chemical reactions Noble gases already have full 8 valence electrons (or 2 for He) so are not chemically reactive Most atoms do not have full valence shells Atoms will gain, lose, or share electrons (form bonds) with other atoms to achieve stability of 8 electrons in valence shell © 2016 Pearson Education, Inc.

Figure 2.5a Chemically inert and reactive elements.
Chemically inert elements Outermost energy level (valence shell) complete 8e 2e 2e He Ne Helium (He) (2p+; 2n0; 2e−) Neon (Ne) (10p+; 10n0; 10e−) © 2016 Pearson Education, Inc.

Figure 2.5b Chemically inert and reactive elements.
Chemically reactive elements Outermost energy level (valence shell) incomplete 4e 1e 2e H C Hydrogen (H) (1p+; 0n0; 1e−) Carbon (C) (6p+; 6n0; 6e−) 1e 6e 8e 2e 2e O Na Oxygen (O) (8p+; 8n0; 8e−) Sodium (Na) (11p+; 12n0; 11e−) © 2016 Pearson Education, Inc.

Types of Chemical Bonds
Three major types of chemical bonds Ionic bonds Covalent bonds Hydrogen bonds © 2016 Pearson Education, Inc.

Types of Chemical Bonds (cont.)
Ionic bonds Ions are atoms that have gained or lost electrons and become charged Number of protons does not equal number of electrons © 2016 Pearson Education, Inc.

Ionic bonds involve the transfer of valence shell electrons from one atom to another, resulting in ions One becomes an anion (negative charge) Atom that gained one or more electrons One becomes a cation (positive charge) Atom that lost one or more electrons Attraction of opposite charges results in an ionic bond © 2016 Pearson Education, Inc.

Figure 2.6ab Formation of an ionic bond.
+  Na Cl Na Cl Sodium atom (Na) (11p+; 12n0; 11e−) Chlorine atom (Cl) (17p+; 18n0; 17e−) Sodium ion (Na+) Chloride ion (Cl−) Sodium chloride (NaCl) Sodium gains stability by losing one electron, and chlorine becomes stable by gaining one electron. After electron transfer, the oppositely charged ions formed attract each other. © 2016 Pearson Education, Inc.

Figure 2.6c Formation of an ionic bond.
Cl− Na+ Large numbers of Na+ and Cl− ions associate to form salt (NaCl) crystals. © 2016 Pearson Education, Inc.

Covalent bonds Covalent bonds are formed by sharing of two or more valence shell electrons between two atoms Sharing of 2 electrons results in a single bond Sharing of 4 electrons is a double bond Sharing of 6 electrons is a triple bond Allows each atom to fill its valence shell at least part of the time Two types of covalent bonds: Polar and nonpolar covalent bonds © 2016 Pearson Education, Inc.

Figure 2.7a Formation of covalent bonds.
Reacting atoms Resulting molecules H H H C H C H or H Structural formula shows single bonds. H H Hydrogen atoms Carbon atom Molecule of methane gas (CH4) Formation of four single covalent bonds: Carbon shares four electron pairs with four hydrogen atoms. © 2016 Pearson Education, Inc.

Figure 2.7b Formation of covalent bonds.
Reacting atoms Resulting molecules O O O O or Structural formula shows double bond. Oxygen atom Oxygen atom Molecule of oxygen gas (O2) Formation of a double covalent bond: Two oxygen atoms share two electron pairs. © 2016 Pearson Education, Inc.

Figure 2.7c Formation of covalent bonds.
Reacting atoms Resulting molecules N N N N or Structural formula shows triple bond. Nitrogen atom Nitrogen atom Molecule of nitrogen gas (N2) Formation of a triple covalent bond: Two nitrogen atoms share three electron pairs. © 2016 Pearson Education, Inc.

Carbon dioxide (CO2) molecules are
Figure 2.8a Carbon dioxide and water molecules have different shapes, as illustrated by molecular models. Carbon dioxide (CO2) molecules are linear and symmetrical. They are nonpolar. © 2016 Pearson Education, Inc.

Polar covalent bonds Unequal sharing of electrons between 2 atoms Results in electrically polar molecules Atoms have different electron-attracting abilities, leading to unequal sharing Atoms with greater electron-attracting ability are electronegative, and those with less are electropositive © 2016 Pearson Education, Inc.

Polar covalent bonds (cont.) H2O is a polar molecule Oxygen is more electronegative, so it exerts a greater pull on shared electrons, giving it a partial negative charge and giving H a partial positive charge Having two different charges is referred to as dipole © 2016 Pearson Education, Inc.

V-shaped water (H2O) molecules have two
Figure 2.8b Carbon dioxide and water molecules have different shapes, as illustrated by molecular models. d− d+ d+ V-shaped water (H2O) molecules have two poles of charge—a slightly more negative oxygen end (d−) and a slightly more positive hydrogen end (d+). © 2016 Pearson Education, Inc.

Ionic bond Polar covalent bond Nonpolar covalent bond Complete
Figure 2.9 Ionic, polar covalent, and nonpolar covalent bonds compared along a continuum. Ionic bond Polar covalent bond Nonpolar covalent bond Complete transfer of electrons Unequal sharing of electrons Equal sharing of electrons Separate ions (charged particles) form Slight negative charge (d−) at one end of molecule, slight positive charge (d+) at other end Charge balanced among atoms Na+ Cl− Sodium chloride Water Carbon dioxide © 2016 Pearson Education, Inc.

Hydrogen bonds Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule Not true bond, more of a weak magnetic attraction Common between dipoles such as water What makes water liquid Also act as intramolecular bonds, holding a large molecule in a three-dimensional shape © 2016 Pearson Education, Inc.

Figure 2.10a Hydrogen bonding between polar water molecules.
(indicated by dotted line) d+ d+ d− d− d− H H O O d+ d+ H H H d+ O H d− The slightly positive ends (d+) of the water molecules become aligned with the slightly negative ends (d−) of other water molecules. © 2016 Pearson Education, Inc.

Figure 2.10b Hydrogen bonding between polar water molecules.
A water strider can walk on a pond because of the high surface tension of water, a result of the combined strength of its hydrogen bonds. © 2016 Pearson Education, Inc.

2.5 Chemical Reactions Chemical Equations
Chemical reactions occur when chemical bonds are formed, rearranged, or broken These reactions can be written in symbolic forms called chemical equations Chemical equations contain: Reactants: substances entering into reaction together Product(s): resulting chemical end products Amounts of reactants and products are shown in balanced equations © 2016 Pearson Education, Inc.

Chemical Equations (cont.)
Compounds are represented as molecular formulas Example: H2O or C6H12O6 Subscript indicates atoms joined by bonds Prefix denotes number of unjoined atoms or molecules © 2016 Pearson Education, Inc.

Chemical Equations (cont.)
Compounds are represented as molecular formulas Example: H2O or C6H12O6 or H2 or CH4 In chemical equations, subscripts indicate how many atoms are joined by bonds, whereas prefix means number of unjoined atoms (example: 4H) Reactants H + H  4H + 1C  Product H2 (Hydrogen gas) CH4 (Methane) © 2016 Pearson Education, Inc.

Types of Chemical Reactions
Three main types of chemical reactions: Synthesis (combination) reactions involve atoms or molecules combining to form larger, more complex molecule Used in anabolic (building) processes A + B  AB © 2016 Pearson Education, Inc.

Figure 2.11a Types of chemical reactions.
Synthesis reactions Smaller particles are bonded together to form larger, more complex molecules. Example Amino acids are joined together to form a protein molecule. Amino acid molecules Protein molecule © 2016 Pearson Education, Inc.

Types of Chemical Reactions (cont.)
Decomposition reactions involve breakdown of a molecule into smaller molecules or its constituent atoms (reverse of synthesis reactions) Involve catabolic (bond-breaking) reactions AB  A + B © 2016 Pearson Education, Inc.

Figure 2.11b Types of chemical reactions.
Decomposition reactions Bonds are broken in larger molecules, resulting in smaller, less complex molecules. Example Glycogen is broken down to release glucose molecules. Glycogen Glucose molecules © 2016 Pearson Education, Inc.

Figure 2.11c Types of chemical reactions.
Exchange reactions Bonds are both made and broken (also called displacement reactions). Example ATP transfers its terminal phosphate group to glucose to form glucose- phosphate. P P P Adenosine triphosphate (ATP) Glucose P P P Adenosine diphosphate (ADP) Glucose- phosphate © 2016 Pearson Education, Inc.

In living systems, these reactions are also referred to as reduction-oxidation or redox reactions Atoms are reduced when they gain electrons and oxidized when they lose electrons Example: C6H12O6 + 6O2 → 6CO2 + 6H2O + ATP In this example, glucose is oxidized, and oxygen molecule is reduced © 2016 Pearson Education, Inc.

Energy Flow in Chemical Reactions
All chemical reactions are either exergonic or endergonic Exergonic reactions result in a net release of energy (give off energy) Products have less potential energy than reactants Catabolic and oxidative reactions Endergonic reactions result in a net absorption of energy (use up energy) Products have more potential energy than reactants Anabolic reactions © 2016 Pearson Education, Inc.

Reversibility of Chemical Reactions
All chemical reactions are theoretically reversible A + B ←→ AB Chemical equilibrium occurs if neither a forward nor a reverse reaction is dominant Many biological reactions are not very reversible Energy requirements to go backward are too high, or products have been removed © 2016 Pearson Education, Inc.

Rate of Chemical Reactions
The speed of chemical reactions can be affected by: Temperature: increased temperatures usually increase rate of reaction Concentration of reactants: increased concentrations usually increase rate Particle size: smaller particles usually increase rate © 2016 Pearson Education, Inc.

Rate of Chemical Reactions
Catalysts Catalysts increase the rate of reaction without being chemically changed or becoming part of the product Enzymes are biological catalysts © 2016 Pearson Education, Inc.

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