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Chemistry Review Atoms and Elements • All types of matter (solids, liquids and gases) are composed of atoms. • A substance that is composed of only one type of atom is called an element. – Elements are the simplest form of matter with unique chemical properties. They are charted on the periodic table based on some of their chemical characteristics. • There are 24 major elements that have various roles in the body. – These include structural, enzymatic, and homeostatic balance. • Compounds, like water, are formed by combining the atoms of different elements together. • Atoms may create various types of chemical bonds. 3 types of bonds include: – Ionic – Covalent – Hydrogen Major Elements of the Human Body • Oxygen (O): Required for energy production during cellular respiration • Carbon (C): Organic back bone for fats, carbohydrates, amino acids and nucleic acids. • Hydrogen (H): Vital for energy( ATP) production • Nitrogen (N): Most abundant in atmosphere, Characteristic element of protein • Phosphorus (P): found in – DNA: Blue print for life – RNA : Vital for protein production – ATP : Cellular energy The Atom • Atoms – identical building blocks for each element • Atomic symbol – one- or two-letter chemical short hand for each element • Atomic number : # of protons in nucleus – periodic table • elements arranged by atomic number • Atomic weight – equal to the mass of the protons and neutrons • Isotope – atoms with same number of protons but a different number of neutrons 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 – valence electrons are located in the outermost shell • interact with other atoms to form bonds – Electrons – have a negative charge and 1/2000 the mass of a proton (0 amu) 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 – except for the first shell which is full with two electrons, atoms interact in a manner to have eight electrons in their valence shell. Planetary Models of Elements p+ represents protons, no represents neutrons Chemically Reactive Elements • Reactive elements do not have their outermost energy level fully occupied by electrons therefore are able to interact with other elements Chemically Inert Elements • Inert elements have their outermost energy level fully occupied by electrons therefore don’t interact with other elements. Types of Chemical Bonds • Ionic: • Covalent • Hydrogen Formation of an Ionic Bond • Ionic bonds form between atoms by the transfer of one or more electrons • Ionic compounds form crystals instead of individual molecules • Example: Na+Cl-(sodium chloride) Formation of an Ionic Bond • A valance electron from Na is transferred to Cl • Cl now has 18e and 17p resulting in a – charge • Na has 10e and 11P resulting in a + charge. Ion Formation • Ions are charged atoms resulting from the gain or loss of electrons. – Anions have gained one or more electrons therefore are negatively charged(-) – Cations have lost one or more electrons giving them a positive charge(+) • Typically occur between elements on opposite sides of the periodic table. Electronegativity and Bond Formation • Elements on opposite ends of the periodic tables have a greater electronegative gradient. Ionic bonds result. • Elements that are closer to each other have smaller electronegative gradient thus form covalent bonds. Covalent bonds • Covalent bonds are formed by the sharing of two or more electrons. – Covalent bonds are classified as Polar or Nonpolar. • When two atoms with similar electronegativities they share their valance electrons. – Nonpolar( neutral charge) bond results. • CO2, O2, N2 • If there is a larger electronegative gradient between the atoms. – a polar covalent bond (charged compound) results. • H 2O Nonpolar Bonds • Electrons shared equally between atoms produce nonpolar bonds. • The negative charged electrons are spaced evenly between the 2 atoms resulting in a neutral charge. Covalent Bonds Double Covalent Bonds Polar Covalent Bonds • Uneven sharing of electrons produces polar bonds • One atom has a greater electronegativity. – This atom will have stronger pull on the shared electrons – The shared electrons spend more time closer to the nucleus of electronegative atom. – The addition of the shared electrons makes the electronegative atom partially negative charged, while the atom with a lower electronegativity becomes partially positively charged – Polar bonds occur between an electronegative atom (mostly O or N) • ex. H2O Inorganic compounds • Do not contain carbon • Water, salts, and many acids and bases • Minerals such as magnesium and calcium. • Salts : NA+CL– – contain cations other than H+ and anions other than OH– • Are electrolytes; they conduct electrical currents and function in various metabolic reactions. – Electrical activity of the nervous system – Vital for bone formation Acid-Base Concentration (pH) • pH scale ranges from 0 to 14. • Acidic solutions have higher [H+] and a lower pH. – Considered proton donors – pH less than 7 • Alkaline (basic) solutions have lower [H+] and a higher pH – Considered proton acceptors – pH greater than 7 • Neutral solutions have equal H+ and OH– concentrations – pH = 7 pH Scale • Acids release H+ and are therefore proton donors HCl H+ + Cl – • Bases release OH– and are proton acceptors NaOH Na+ + OH– Buffers • The body has many mechanisms devoted to resist abrupt and large swings in the pH of body fluids. • These systems allow pH to remain relatively constant . – Approximately 7.4 (slightly basic) – Maintaining a stable pH is critical for creating an environment necessary for metabolic reactions. Chemical Reactions • Chemical reactions in the body act by forming, breaking or rearranging bonds. • Chemical equations contain: – Relative amounts of reactants (starting chemicals) and products (finishing chemicals) – Number and type of reacting substances, and products produced Synthesis and Decomposition Reactions Oxidation-Reduction (Redox) Reactions • Reactants losing electrons are electron donors and are oxidized • Reactants taking up electrons are electron acceptors and become reduced • Na + Cl → Na+ + Cl– Na is oxidized and Cl has been reduced • LEO THE LION SAYS GER Forms of Energy • Chemical – stored in the bonds of chemical substances – Energy from food • Electrical – results from the movement of charged particles – Household Appliances run on it • Mechanical – directly involved in moving matter – Machines such as cranes or bull dozers • Radiant or electromagnetic – energy traveling in waves – visible light, ultraviolet light, and X rays First Law of Thermodynamics – Energy cannot be created or destroyed, but only change form. • During each conversion, some of the energy dissipates into the environment as heat. –Heat is defined as the measure of the random motion of molecules. – the second law states that "energy systems have a tendency to increase their entropy" Energy • The capacity to do work (put matter into motion) • Types of energy – Kinetic – energy in action. Ball rolling down a hill. – Potential – energy of position; stored (inactive) energy. Ball sitting on top of hill. Fig. 8.2 (TEArt) Potential energy Energy - the capacity to do work –kinetic - energy of motion –potential - stored energy • Kinetic energy Factors Influencing Rate of Chemical Reactions • Temperature – chemical reactions proceed quicker at higher temperatures • Particle size – the smaller the particle the faster the chemical reaction • Concentration – higher reacting particle concentrations produce faster reactions • Catalysts – increase the rate of a reaction without being chemically changed • Enzymes – biological catalysts