Chapter 2 PowerPoint Amerman PDF

ERIN C. AMERMAN FLORIDA STATE COLLEGE AT JACKSONVILLE Lecture Presentation by Wallace State Community College © 2016 Pearson Education, Inc.  MATTER Matter – anything that has mass and occupies space;...

ERIN C. AMERMAN FLORIDA STATE COLLEGE AT JACKSONVILLE Lecture Presentation by Wallace State Community College © 2016 Pearson Education, Inc.  MATTER Matter – anything that has mass and occupies space; can exist in three states: solid, liquid, or gas Chemistry – study of matter and its interactions Atom – smallest unit of matter that retains original properties © 2016 Pearson Education, Inc.  ATOMS AND ATOMIC STRUCTURE Atoms are made up of even smaller structures called subatomic particles Subatomic particles exist in 3 forms:  Protons (p+) – found in central core of atom (atomic nucleus); positively charged  Neutrons (n0) – found in atomic nucleus; slightly larger than protons; no charge.  Electrons (e-) – found outside atomic nucleus; negatively charged Atoms are electrically neutral – they have no charge; number of protons and electrons are equal, cancelling each other’s charge; number of neutrons does not have to equal number of protons Figure 2.1 Structure of a representative atom. © 2016 Pearson Education, Inc.  ATOMS AND ATOMIC STRUCTURE Electron shells – regions surrounding atomic nucleus where electrons exist; each can hold a certain number of electrons:  1st shell (closest to nucleus) can hold 2 electrons  2nd shell can hold 8 electrons © 2016 Pearson Education, Inc.  ELEMENTS IN THE PERIODIC TABLE AND THE HUMAN BODY Number of protons that an atom has in its nucleus is its atomic number Atomic number defines every element:  Element – substance that cannot be broken down into simpler substance by chemical means  Each element is made of atoms with same number of protons © 2016 Pearson Education, Inc.  ELEMENTS IN THE PERIODIC TABLE AND THE HUMAN BODY The human body is made up of four major elements:  Hydrogen  Oxygen  Carbon  Nitrogen Also 7 mineral elements and 13 trace elements © 2016 Pearson Education, Inc.  ISOTOPES AND RADIOACTIVITY Mass number – equal to sum of all protons and neutrons found in atomic nucleus Isotope – atom with same atomic number (same number of protons), but different mass number (different number of neutrons) Radioisotopes – unstable isotopes; high energy or radiation released by radioactive decay; allows isotope to assume a more stable form © 2016 Pearson Education, Inc.  NUCLEAR MEDICINE Common applications of radioisotopes: Cancer radiation therapy – radiation damages structure of cancer cells; interferes with functions Radiotracers – injected into patient and detected by camera; image analyzed by computer; shows size, shape, and activity of organs and cells Treatment of thyroid disorders – high doses of iodine-131 treat overactive or cancerous thyroid tissue; radioisotope accumulates and damages cells © 2016 Pearson Education, Inc.  MATTER COMBINED Matter can be combined physically to form a mixture – atoms of two or more elements physically intermixed without changing chemical nature of atoms themselves There are 3 basic types of mixtures: suspensions, colloids, and solutions (Figure 2.3) © 2016 Pearson Education, Inc.  MIXTURES Suspension – mixture containing two or more components with large, unevenly distributed particles; will settle out when left undisturbed Colloids – two or more components with small, evenly distributed particles; will not settle out Figure 2.3a The three types of mixtures. © 2016 Pearson Education, Inc.  MIXTURES Solutions – two or more components with extremely small, evenly distributed particles; will not settle out; contain a solute dissolved in a solvent:  Solute – substance that is dissolved  Solvent – substance that dissolves solute Figure 2.3c The three types of mixtures. © 2016 Pearson Education, Inc.  CHEMICAL BONDS Matter can be combined chemically when atoms are combined by chemical bonds. A chemical bond is not a physical structure but rather an energy relationship or attractive force between atoms  Molecule – formed by chemical bonding between two or more atoms of same element  Compound – formed when two or more atoms from different elements combine by chemical bonding CH4 © 2016 Pearson Education, Inc.  CHEMICAL BONDS Macromolecules – very large molecules composed of many atoms Molecular formulas – represent molecules symbolically with letters and numbers; show kinds and numbers of atoms in a molecule CH4 O2 N2 Table 2.1 Electron Sharing in Covalent Bonds. © 2016 Pearson Education, Inc.  CHEMICAL BONDS Chemical bonds are formed when valence electrons (in outermost valence shell) of atoms interact Valence electrons determine how an atom interacts with other atoms and whether it will form bonds with a specific atom  The octet rule states that an atom is most stable when it has 8 electrons in its valence shell (as in CO2). It is considered nonreactive or inert if the outer ring is filled. © 2016 Pearson Education, Inc.  IONS AND IONIC BONDS Ionic bond – formed when electrons are transferred from a metal atom to a nonmetal atom; results in formation of ions: cations and anions (Figure 2.4)  Cation – positively charged ion; forms when metal loses one or more electrons  Anion – negatively charged ion; forms when nonmetal gains one or more electrons The attraction between opposite charges bonds ions to one another forming a compound called a salt © 2016 Pearson Education, Inc.  IONIC BONDS Figure 2.4 Formation of an ionic bond. © 2016 Pearson Education, Inc.  COVALENT BONDS Covalent bonds – strongest bond; form when two or more nonmetals share electrons (Figures 2.5, 2.6; Table 2.1) Two atoms can share one (single bond), two (double bond), or three (triple bond) electron pairs: Table 2.1 Electron Sharing in Covalent Bonds. © 2016 Pearson Education, Inc.  COVALENT BONDS All elements have protons that attract electrons; property known as electronegativity: The more electronegative an element the more strongly it attracts electrons, pulling them away from less electronegative elements © 2016 Pearson Education, Inc.  NONPOLAR COVALENT BONDS Nonpolar covalent bonds result when two nonmetals in a molecule with similar or identical electronegativities pull with equal force; therefore share electrons equally (Figure 2.6a) Nonpolar molecules occur in 3 situations: Atoms sharing electrons are same element Arrangement of atoms makes one atom unable to pull more strongly than another atom (as in CO2) Bond is between carbon and hydrogen © 2016 Pearson Education, Inc.  POLAR COVALENT BONDS Polar covalent bonds form polar molecules when nonmetals with different electronegativities interact resulting in an unequal sharing of electrons (Figure 2.6b)  Atom with higher electronegativity becomes partially negative (δ) as it pulls shared electrons close to itself  Atom with lower electronegativity becomes partially positive (δ+) as shared electrons are pulled toward other atom Polar molecules with partially positive and partially negative ends are known as dipoles © 2016 Pearson Education, Inc.  POLAR COVALENT BONDS Figure 2.6b Nonpolar vs. polar covalent bonds. © 2016 Pearson Education, Inc.  HYDROGEN BONDS Hydrogen bonds – weak attractions between partially positive end of one dipole and partially negative end of another dipole Hydrogen bonds are responsible for a key property of water—surface tension Figure 2.7a Hydrogen bonding and surface tension between water molecules. © 2016 Pearson Education, Inc.  HYDROGEN BONDS Polar water molecules are more strongly attracted to one another than they are to nonpolar air molecules at surface Figure 2.7b Hydrogen bonding and surface tension between water molecules. © 2016 Pearson Education, Inc.  CHEMICAL NOTATION A chemical reaction has occurred every time a chemical bond is formed, broken, or rearranged, or when electrons are transferred between two or more atoms (or molecules) Chemical notation – series of symbols and abbreviations used to demonstrate what occurs in a reaction; the chemical equation has two parts:  Reactants on left side of equation are starting ingredients; will undergo reaction  Products on right side of equation are results of chemical reaction © 2016 Pearson Education, Inc.  CHEMICAL NOTATION Reversible reactions can proceed in either direction as denoted by two arrows that run in opposite directions (as below) Irreversible reactions proceed from left to right as denoted by a single arrow CO2 + H2O H2CO3 Reactants (carbon dioxide + water) Product (carbonic acid) © 2016 Pearson Education, Inc.  ENERGY AND CHEMICAL REACTIONS Energy is defined as capacity to do work or put matter into motion or fuel chemical reactions; two general forms of energy:  Potential energy is stored; can be released to do work at some later time  Kinetic energy is potential energy that has been released or set in motion to perform work; all atoms have kinetic energy as they are in constant motion; the faster they move the greater that energy © 2016 Pearson Education, Inc.  ENERGY AND CHEMICAL REACTIONS Energy is found in 3 forms in the human body; chemical, electrical, and mechanical, each of which may be potential or kinetic depending on location or process Energy, inherent in all chemical bonds, must be invested any time a chemical reaction occurs: Endergonic reactions require input of energy from another source; products contain more energy than reactants because energy was invested so reaction could proceed Exergonic reactions release excess energy so products have less energy than reactants © 2016 Pearson Education, Inc.  HOMEOSTASIS AND TYPES OF CHEMICAL REACTIONS Three fundamental processes occur in the body to maintain homeostasis (breaking down molecules, converting the energy in food to usable form, and building new molecules); carried out by three basic types of chemical reactions: 1. Catabolic reactions (decomposition reactions) 2. Exchange reactions 3. Anabolic reactions (synthesis reactions) © 2016 Pearson Education, Inc.  HOMEOSTASIS AND TYPES OF CHEMICAL REACTIONS Catabolic reactions (decomposition reactions) – when a large substance is broken down into smaller substances General chemical notation for reaction is AB A + B Usually exergonic because chemical bonds are broken © 2016 Pearson Education, Inc.  HOMEOSTASIS AND TYPES OF CHEMICAL REACTIONS Exchange reactions occur when one or more atoms from reactants are exchanged for one another General chemical notation for reaction is AB + CD AD + BC HCL + NaOH H2O + NaCL © 2016 Pearson Education, Inc.  HOMEOSTASIS AND TYPES OF CHEMICAL REACTIONS Oxidation-reduction reactions (redox reactions) – special kind of exchange reaction; occur when electrons and energy are exchanged instead of atoms  Reactant that loses electrons is oxidized  Reactant that gains electrons is reduced Redox reactions are usually exergonic reactions capable of releasing large amounts of energy © 2016 Pearson Education, Inc.  HOMEOSTASIS AND TYPES OF CHEMICAL REACTIONS Anabolic reactions (synthesis reactions) occur when small simple subunits and united by chemical bonds to make large more complex substances General chemical notation for reaction is A + B AB These reactions are endergonic; fueled by chemical energy © 2016 Pearson Education, Inc.  REACTION RATES AND ENZYMES For a reaction to occur atoms must collide with enough energy overcome the repulsion of their electrons This energy required for all chemical reactions is called the activation energy (Ea) Figure 2.8 Activation energy. © 2016 Pearson Education, Inc.  REACTION RATES AND ENZYMES Analogy can be applied to chemical reactions – activation energy must be supplied so that reactants reach their transition states (i.e., get to the top of the energy “hill”) in order to react and form products (i.e., roll down the hill) Figure 2.8 Activation energy. © 2016 Pearson Education, Inc.  REACTION RATES AND ENZYMES The following factors increase reaction rate by reducing activation energy or increasing likelihood of strong collisions between reactants:  Concentration  Temperature  Reactant properties  Presence or absence of a catalyst © 2016 Pearson Education, Inc.  REACTION RATES AND ENZYMES When reactant concentration increases, more reactant particles are present, increasing chance of successful collisions between reactants Raising the temperature of the reactants increases kinetic energy of their atoms leading to more forceful and effective collisions between reactants Both particle size and phase (solid, liquid, or gas) influence reaction rates:  Smaller particles move faster with more energy than larger particles  Reactant particles in the gaseous phase have higher kinetic energy than those in either solid or liquid ph © 2016 Pearson Education, Inc.  ENZYME ACTIVITY In Groups © 2016 Pearson Education, Inc.  REACTION RATES AND ENZYMES Figure 2.9 The effect of enzymes on activation energy. © 2016 Pearson Education, Inc.  ENZYME DEFICIENCIES Examples of common enzyme deficiencies: Tay-Sachs Disease – deficiency of hexosaminidase; gangliosides accumulate around neurons of brain; death usually by age 3 Severe Combined Immunodeficiency Syndrome (SCIDs) – may be due to adenosine deaminase deficiency; nearly complete absence of immune system; affected patients must live in sterile “bubble” Phenylketonuria – deficiency of phenylalanine hydroxylase; converts phenylalanine into tyrosine; resulting seizures and mental retardation can be prevented by dietary modification © 2016 Pearson Education, Inc.  BIOCHEMISTRY Biochemistry – the chemistry of life Inorganic compounds generally do not contain carbon bonded to hydrogen; include water, acids, bases, and salts Organic compounds – those that do contain carbon bonded to hydrogen © 2016 Pearson Education, Inc.  WATER Water (H2O) makes up 60–80% of mass of human body and has several key properties vital to our existence (Figure 2.11): High heat capacity – able to absorb heat without significantly changing temperature itself Carries heat with it when it evaporates (when changing from liquid to gas) Cushions and protects body structures because of relatively high density Acts as a lubricant between two adjacent surfaces (reduces friction) Water serves as body’s primary solvent; often called the universal solvent because so many solutes will dissolve in it entirely or to some degree (Figure 2.11) © 2016 Pearson Education, Inc.  WATER Water is only able to dissolve hydrophilic solutes (those with fully or partially charged ends); “like dissolves like”, so water dissolves ionic and polar covalent solutes Figure 2.11a, b The behavior of hydrophilic and hydrophobic molecules in water. © 2016 Pearson Education, Inc.  WATER Solutes that do not have full or partially charged ends are hydrophobic; do not dissolve in water; includes uncharged nonpolar covalent molecules such as oils and fats Figure 2.11c The behavior of hydrophilic and hydrophobic molecules in water. © 2016 Pearson Education, Inc.  ACIDS AND BASES Figure 2.12 The behavior of acids and bases in water. © 2016 Pearson Education, Inc.  ACIDS AND BASES Figure 2.13 The pH Scale. © 2016 Pearson Education, Inc.  SALTS AND ELECTROLYTES Salt – any metal cation and nonmetal anion held together by ionic bonds Salts can dissolve in water to form cations and anions called electrolytes which are capable of conducting electrical current Figure 2.4 and Figure 2.11a © 2016 Pearson Education, Inc. MONOMERS AND POLYMERS Each type of organic compound in body (carbohydrate, lipid, protein, or nucleic acid) consists of polymers built from monomer subunits: Monomers are single subunits that can be combined to build larger structures called polymers by dehydration synthesis (anabolic reaction that links monomers together and makes a molecule of water in process) Hydrolysis is a catabolic reaction that uses water to break up polymers into smaller subunits © 2016 Pearson Education, Inc. DO BUILDING POLYMER ACTIVITY- GLUCOSE © 2016 Pearson Education, Inc.  CARBOHYDRATES Carbohydrates, composed of carbon, hydrogen, and oxygen, function primarily as fuel; some limited structural roles  Monosaccharides – consist of 3 to 7 carbons; monomers from which all carbohydrates are made; glucose, fructose, galactose, ribose, and dexoyribose are most abundant monosaccharides (Figure 2.14)  Disaccharides are formed by union of two monosaccharides by dehydration synthesis © 2016 Pearson Education, Inc.  CARBOHYDRATES Polysaccharides consist of many monosaccharides joined to one another by dehydration synthesis reactions (Figure 2.16) Glycogen is the storage polymer of glucose; mostly in skeletal muscle and liver cells Starch – dietary polysaccharide from plant material © 2016 Pearson Education, Inc.  LIPIDS Lipids – group of nonpolar hydrophobic molecules composed primarily of carbon and hydrogen; include fats and oils © 2016 Pearson Education, Inc.  LIPIDS Saturated fatty acids – solid at room temperature; have no double bonds between carbon atoms so carbons are “saturated” with maximum number of hydrogen atoms Figure 2.17a Lipids: structure of fatty acids. © 2016 Pearson Education, Inc.  LIPIDS  Monounsaturated fatty acids – generally liquid at room temperature; have one double bond between two carbons in hydrocarbon chain  Polyunsaturated fatty acids – liquid at room temperature; have two or more double bonds between carbons in hydrocarbon chain Figure 2.17b Lipids: structure of fatty acids. © 2016 Pearson Education, Inc.  LIPIDS Triglyceride – three fatty acids linked by dehydration synthesis to a modified 3-carbon carbohydrate, glycerol; storage polymer for fatty acids (also called a neutral fat) Figure 2.18 Lipids: structure and formation of triglycerides. © 2016 Pearson Education, Inc.  LIPIDS Phospholipids – composed of a glycerol backbone, two fatty acid “tails” and one phosphate “head” in place of third fatty acid Makes up the cell membrane Table 2.3 Organic Molecules. © 2016 Pearson Education, Inc.  LIPIDS Steroids – nonpolar and share a four- ring hydrocarbon structure called the steroid nucleus Cholesterol – steroid that forms basis for all other steroids Figure 2.20 Lipids: structure of steroids and Table 2.3 Organic Molecules. © 2016 Pearson Education, Inc.  PROTEINS Proteins are macromolecules that:  Function as enzymes  Play structural roles  Are involved in movement  Function in the body’s defenses  Can be used as fuel © 2016 Pearson Education, Inc.  PROTEINS Twenty different amino acids (monomers of all proteins); can be linked by peptide bonds into polypeptides Figure 2.21a, b Proteins: structure of amino acids. © 2016 Pearson Education, Inc.  PROTEINS Peptides – formed from two or more amino acids linked together by peptide bonds through dehydration synthesis: Figure 2.22 Proteins: formation and breakdown of dipeptides. © 2016 Pearson Education, Inc.  PROTEINS Proteins consist of one or more polypeptide chains folded into distinct structures which must be maintained to be functional; example of Structure-Function Core Principle Complex structure of a complete protein is divided into four levels: Figure 2.23a Levels of protein structure. © 2016 Pearson Education, Inc.  PROTEINS Protein denaturation – process of destroying a protein’s shape by heat, pH changes, or exposure to chemicals Disrupts hydrogen bonding and ionic interactions that stabilize structure and function. © 2016 Pearson Education, Inc. NUCLEOTIDES AND NUCLEIC ACIDS Nucleotides – monomers of nucleic acids; named because of abundance in nuclei of cells; make up genetic material Nucleotide structure:  Nitrogenous base with a hydrocarbon ring structure  Five-carbon pentose sugar, ribose or dexoyribose  Phosphate group Figure 2.24a Structure of nucleotides. © 2016 Pearson Education, Inc. NUCLEOTIDES AND NUCLEIC ACIDS Two types of nitrogenous bases: purines and pyrimidines Purines – double-ringed molecule; adenine (A) and guanine (G) Pyrimidines – single-ringed molecule; cytosine (C), uracil (U) and thymine (T) Figure 2.24 Structure of nucleotides. © 2016 Pearson Education, Inc. NUCLEOTIDES AND NUCLEIC ACIDS Adenosine triphosphate (ATP) Adenine attached to ribose and three phosphate groups; main source of chemical energy in body Synthesized from adenosine diphosphate (ADP) and a phosphate group (Pi) using energy from oxidation of fuels (like glucose) Figure 2.25a Nucleotides: structure and formation of ATP. © 2016 Pearson Education, Inc. NUCLEOTIDES AND NUCLEIC ACIDS Adenosine triphosphate (continued): Potential energy in this “high-energy” bond can be released as kinetic energy to do work Production of large quantities of ATP requires oxygen; why we breathe air Figure 2.25b Nucleotides: structure and formation of ATP. © 2016 Pearson Education, Inc. NUCLEOTIDES AND NUCLEIC ACIDS DNA, an extremely large molecule found in nuclei of cells; composed of two long chains that twist around each other to form a double helix DNA contains genes – provide recipe or code for protein synthesis – process of making every protein Figure 2.26a Structure of nucleic acids and Table 2.3 Organic Molecules. © 2016 Pearson Education, Inc. NUCLEOTIDES AND NUCLEIC ACIDS Other structural features of DNA include: DNA contains:  Pentose sugar deoxyribose (lacks oxygen-containing group of ribose) forms backbone of strand; alternates with phosphate group  Bases: adenine, guanine, cytosine, and thymine Figure 2.26a Structure of nucleic acids and Table 2.3 Organic Molecules. © 2016 Pearson Education, Inc. NUCLEOTIDES AND NUCLEIC ACIDS Other structural features of DNA include: Double helix strands – held together by hydrogen bonding between the bases of each strand DNA exhibits complementary base pairing; purine A always pairs with pyrimidine T and purine G always pairs with pyrimidine C A = T (where = denotes 2 hydrogen bonds) and C  G (where  denotes 3 hydrogen bonds) © 2016 Pearson Education, Inc. NUCLEOTIDES AND NUCLEIC ACIDS RNA – single strand of nucleotides; can move between nucleus of cell and cytosol; critical to making proteins RNA contains the pentose sugar ribose RNA contains uracil instead of thymine; still pairs with adenine, (A = U) Figure 2.26b Structure of the nucleic acids DNA and RNA. © 2016 Pearson Education, Inc.  NUCLEOTIDES AND NUCLEIC ACIDS Figure 2.26 Structure of the nucleic acids DNA and RNA. © 2016 Pearson Education, Inc. DO MACROMOLECULE PRACTICE © 2016 Pearson Education, Inc.