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Biological Sciences
Kenneth Welch

Chemistry 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) 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 Major Elements of the Human Body • Oxygen (O) • Carbon (C) • Hydrogen (H) • Nitrogen (N) 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 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 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 – 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 • Suspensions are heterogeneous mixtures with visible solutes that tend to settle out 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 – atoms react in a manner to have 8 electrons in their valence shell Chemically Inert and Reactive Elements • Inert elements have their outermost energy level fully occupied by electrons Chemically Inert and Reactive Elements • Reactive elements do not have their outermost energy level fully occupied by electrons 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 • Opposite charges on anions and cations hold them close together, forming ionic bonds Formation of an Ionic Bond • Ionic compounds form crystals instead of individual molecules • Example: NaCl (sodium chloride) Formation of an Ionic Bond Covalent Bonds • Electrons are shared by two atoms • Electron sharing produces molecules Polar and Nonpolar Molecules • Electrons shared equally between atoms produce nonpolar molecules • Unequal sharing of electrons produces polar molecules • Atoms with 6 or 7 valence shell electrons are electronegative • Atoms with 1 or 2 valence shell electrons are electropositive 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 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 H + H  H 2 (reactants) (product) 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 Energy Flow in Chemical Reactions • Exergonic reactions – reactions that release energy • Endergonic reactions – reactions whose products contain more potential energy than did its reactants Reversibility of Chemical Reactions • All chemical reactions are theoretically reversible
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