Chemistry Comes Alive: Part A 2 Chemistry Comes Alive: Part A

Anything that has mass and occupies space States of matter: Solid—definite shape and volume Liquid—definite volume, changeable shape Gas—changeable shape and volume

Capacity to do work or put matter into motion Types of energy: Kinetic—energy in action Potential—stored (inactive) energy PLAY Animation: Energy Concepts

Forms of Energy Chemical 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 —exhibits wavelike properties (i.e., visible light, ultraviolet light, and X-rays)

Energy Form Conversions Energy may be converted from one form to another Conversion is inefficient because some energy is “lost” as heat

Composition of Matter Elements Cannot be broken down by ordinary chemical means Each has unique properties: Physical properties Are detectable with our senses, or are measurable Chemical properties How atoms interact (bond) with one another

Atomic symbol: one- or two-letter chemical shorthand for each element Composition of Matter Atoms Unique building blocks for each element Atomic symbol: one- or two-letter chemical shorthand for each element

Major Elements of the Human Body Oxygen (O) Carbon (C) Hydrogen (H) Nitrogen (N) About 96% of body mass

Lesser Elements of the Human Body About 3.9% of body mass: Calcium (Ca), phosphorus (P), potassium (K), sulfur (S), sodium (Na), chlorine (Cl), magnesium (Mg), iodine (I), and iron (Fe)

Trace Elements of the Human Body < 0.01% of body mass: Part of enzymes, e.g., chromium (Cr), manganese (Mn), and zinc (Zn)

Atomic Structure – Review Write this down if needed! Determined by numbers of subatomic particles Nucleus consists of neutrons and protons

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Atomic Structure Neutrons Protons No charge Mass = 1 atomic mass unit (amu) Protons Positive charge Mass = 1 amu

Atomic Structure Electrons Orbit nucleus Equal in number to protons in atom Negative charge 1/2000 the mass of a proton (0 amu)

Nucleus Nucleus Helium atom Helium atom 2 protons (p+) 2 neutrons (n0) 2 electrons (e–) 2 protons (p+) 2 neutrons (n0) 2 electrons (e–) (a) Planetary model (b) Orbital model Proton Neutron Electron Electron cloud Figure 2.1

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Identifying Elements – do not write this Atoms of different elements contain different numbers of subatomic particles Compare hydrogen, helium and lithium (next slide)

Hydrogen (H) (1p+; 0n0; 1e–) Helium (He) (2p+; 2n0; 2e–) Lithium (Li) Proton Neutron Electron Hydrogen (H) (1p+; 0n0; 1e–) Helium (He) (2p+; 2n0; 2e–) Lithium (Li) (3p+; 4n0; 3e–) Figure 2.2

Identifying Elements Atomic number = number of protons in nucleus

Mass number = mass of the protons and neutrons Identifying Elements Mass number = mass of the protons and neutrons Mass numbers of atoms of an element are not all identical Isotopes are structural variations of elements that differ in the number of neutrons they contain

Identifying Elements Atomic weight = average of mass numbers of all isotopes

Valuable tools for biological research and medicine Radioisotopes Valuable tools for biological research and medicine Cause damage to living tissue: Useful against localized cancers Radon from uranium decay causes lung cancer

Molecules and Compounds Most atoms combine chemically with other atoms to form molecules and compounds Molecule—two or more atoms bonded together (e.g., H2) Compound—two or more different kinds of atoms bonded together (e.g., C6H12O6)

Most matter exists as mixtures Two or more components physically intermixed Three types of mixtures Solutions Colloids Suspensions

Usually transparent, e.g., atmospheric air or seawater Solutions Homogeneous mixtures Usually transparent, e.g., atmospheric air or seawater Solvent Present in greatest amount, usually a liquid Solute(s) Present in smaller amounts

Colloids and Suspensions Colloids (emulsions) Heterogeneous translucent mixtures, e.g., cytosol Large solute particles that do not settle out Undergo sol-gel transformations Suspensions: Heterogeneous mixtures, e.g., blood Large visible solutes tend to settle out

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 Mineral water Example Gelatin Example Blood Figure 2.4

Mixtures vs. Compounds Mixtures Compounds No chemical bonding between components Can be separated physically, such as by straining or filtering Heterogeneous or homogeneous Compounds Can be separated only by breaking bonds All are homogeneous

Chemical Bonds – review only Electrons occupy up to seven electron shells (energy levels) around nucleus Octet rule: Except for the first shell which is full with two electrons, atoms interact in a manner to have eight electrons in their outermost energy level (valence shell)

Chemically Reactive Elements – review only Outermost energy level not fully occupied by electrons Tend to gain, lose, or share electrons (form bonds) with other atoms to achieve stability

(b) Chemically reactive elements Outermost energy level (valence shell) incomplete 4e 1e 2e Hydrogen (H) (1p+; 0n0; 1e–) Carbon (C) (6p+; 6n0; 6e–) 1e 6e 8e 2e 2e Oxygen (O) (8p+; 8n0; 8e–) Sodium (Na) (11p+; 12n0; 11e–) Figure 2.5b

Types of Chemical Bonds Ionic Covalent Hydrogen

Ions are formed by transfer of valence shell electrons between atoms Ionic Bonds Ions are formed by transfer of valence shell electrons between atoms Anions (– charge) have gained one or more electrons Cations (+ charge) have lost one or more electrons Attraction of opposite charges results in an ionic bond

+ – Figure 2.6a-b Sodium atom (Na) (11p+; 12n0; 11e–) Chlorine atom (Cl) (17p+; 18n0; 17e–) Sodium ion (Na+) Chloride ion (Cl–) Sodium chloride (NaCl) (a) Sodium gains stability by losing one electron, and chlorine becomes stable by gaining one electron. (b) After electron transfer, the oppositely charged ions formed attract each other. Figure 2.6a-b

Formation of an Ionic Bond Ionic compounds form crystals instead of individual molecules NaCl (sodium chloride)

(c) Large numbers of Na+ and Cl– ions CI– Na+ (c) Large numbers of Na+ and Cl– ions associate to form salt (NaCl) crystals. Figure 2.6c

Covalent Bonds Formed by sharing of two or more valence shell electrons Allows each atom to fill its valence shell at least part of the time

+ Reacting atoms Resulting molecules or Structural formula shows single bonds. Hydrogen atoms Carbon atom Molecule of methane gas (CH4) (a) Formation of four single covalent bonds: carbon shares four electron pairs with four hydrogen atoms. Figure 2.7a

+ Reacting atoms Resulting molecules or Structural formula shows double bond. Oxygen atom Oxygen atom Molecule of oxygen gas (O2) (b) Formation of a double covalent bond: Two oxygen atoms share two electron pairs. Figure 2.7b

+ Reacting atoms Resulting molecules or Structural formula shows triple bond. Nitrogen atom Nitrogen atom Molecule of nitrogen gas (N2) (c) Formation of a triple covalent bond: Two nitrogen atoms share three electron pairs. Figure 2.7c

Sharing of electrons may be equal or unequal Covalent Bonds Sharing of electrons may be equal or unequal Equal sharing produces electrically balanced nonpolar molecules CO2

Figure 2.8a

Covalent Bonds Unequal sharing by atoms with different electron-attracting abilities produces polar molecules H2O Atoms with six or seven valence shell electrons are electronegative, e.g., oxygen Atoms with one or two valence shell electrons are electropositive, e.g., sodium

Figure 2.8b

Figure 2.9

Hydrogen Bonds Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule Common between dipoles such as water Also act as intramolecular bonds, holding a large molecule in a three-dimensional shape PLAY Animation: Hydrogen Bonds

Hydrogen bond (indicated by dotted line) + – Hydrogen bond (indicated by dotted line) + + – – – + + + – (a) The slightly positive ends (+) of the water molecules become aligned with the slightly negative ends (–) of other water molecules. Figure 2.10a

(b) 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. Figure 2.10b

Occur when chemical bonds are formed, rearranged, or broken Chemical Reactions Occur when chemical bonds are formed, rearranged, or broken Represented as chemical equations Chemical equations contain: Molecular formula for each reactant and product Relative amounts of reactants and products, which should balance

Examples of Chemical Equations H + H  H2 (hydrogen gas) 4H + C  CH4 (methane) (reactants) (product)

Patterns of Chemical Reactions Synthesis (combination) reactions A + B  AB Decomposition reactions AB  A + B Exchange reactions AB + C  AC + B

Synthesis Reactions A + B  AB Always involve bond formation Anabolic ( creation of a product, the building up of organs and tissues, requires energy! ATP) Anabolism is powered by Catabolism!

(a) 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 Figure 2.11a

Decomposition Reactions AB  A + B Reverse synthesis reactions Involve breaking of bonds Catabolic (set of metabolic pathways that break down molecules and release energy)

(b) Decomposition reactions Bonds are broken in larger molecules, resulting in smaller, less complex molecules. Example Glycogen is broken down to release glucose units. Glycogen Glucose molecules Figure 2.11b

Exchange Reactions AB + C  AC + B Also called displacement reactions Bonds are both made and broken

+ + (c) 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. + Glucose Adenosine triphosphate (ATP) + Glucose phosphate Adenosine diphosphate (ADP) Figure 2.11c

Oxidation-Reduction (Redox) Reactions Decomposition reactions: Reactions in which fuel is broken down for energy Also called exchange reactions because electrons are exchanged or shared differently Electron donors lose electrons and are oxidized Electron acceptors receive electrons and become reduced

All chemical reactions are either exergonic or endergonic Exergonic reactions —release energy Catabolic reactions Endergonic reactions —products contain more potential energy than did reactants Anabolic reactions

Chemical Reactions – review only All chemical reactions are theoretically reversible A + B  AB AB  A + B Chemical equilibrium occurs if neither a forward nor reverse reaction is dominant Many biological reactions are essentially irreversible due to Energy requirements Removal of products

Rate of Chemical Reactions – review only Rate of reaction is influenced by:  temperature   rate  particle size   rate  concentration of reactant   rate Catalysts:  rate without being chemically changed Enzymes are biological catalysts