Chemistry Comes Alive Part A

1. Chemistry Comes Alive Part A 2

2. 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

3. Energy • The capacity to do work (put matter into motion) • Types of energy • Kinetic – energy in action • Potential – energy of position; stored (inactive) energy

4. 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)

5. Energy Form Conversions • Energy is easily converted from one form to another • During conversion, some energy is “lost” as heat

6. 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

7. Properties of Elements • Each element has unique physical and chemical properties • Physical properties – those detected with our senses • Chemical properties – pertain to the way atoms interact with one another

8. Major Elements of the Human Body • Oxygen (O) • Carbon (C) • Hydrogen (H) • Nitrogen (N) • 96% of body matter

9. 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

10. 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)

11. 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

12. Models of the Atom Figure 2.1

13. 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

14. Identification of Elements Figure 2.2

15. Identification of Elements Figure 2.3

16. Molecules and Compounds • Molecule – two or more atoms held together by chemical bonds • Compound – two or more different kinds of atoms chemically bonded together

17. Mixtures and Solutions • 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

18. 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

19. Colloids and Suspensions • Colloids, or emulsions, are heterogeneous mixtures whose solutes do not settle out • Example: Jello and Cytosol • Suspensions are heterogeneous mixtures with visible solutes that tend to settle out • Example: Blood

20. 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

21. 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

22. Chemically Inert Elements • Inert elements have their outermost energy level fully occupied by electrons Figure 2.4a

23. Chemically Reactive Elements • Reactive elements do not have their outermost energy level fully occupied by electrons Figure 2.4b

24. Types of Chemical Bonds • Ionic • Covalent • Hydrogen

25. 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

26. 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: NaCl (sodium chloride)

27. Formation of an Ionic Bond Figure 2.5a

28. Formation of an Ionic Bond Figure 2.5b

29. Covalent Bonds • Covalent bonds are formed by the sharing of two or more electrons • Electron sharing produces molecules

30. Single Covalent Bonds Figure 2.6a

31. Double Covalent Bonds Figure 2.6b

32. Triple Covalent Bonds Figure 2.6c

33. Polar and Nonpolar Molecules • Electrons shared equally between atoms produce nonpolar molecules • Unequal sharing of electrons produces polar molecules • Atoms with six or seven valence shell electrons are electronegative • Atoms with one or two valence shell electrons are electropositive

34. Comparison of Ionic, Polar Covalent, and Nonpolar Covalent Bonds Figure 2.8

35. 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

36. Hydrogen Bonds Figure 2.9

37. 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

38. Examples of Chemical Reactions

39. 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

40. Oxidation-Reduction (Redox) Reactions • Reactants losing electrons are electron donors and are oxidized • Reactants taking up electrons are electron acceptors and become reduced • Therefore, both decomposition and electron exchange occur.

41. Energy Flow in Chemical Reactions • Exergonic reactions – reactions that release energy • Usually when a bond is broken. • Endergonic reactions – reactions whose products contain more potential energy than did its reactants

42. Reversibility in Chemical Reactions • All chemical reactions are theoretically reversible A + B  AB AB  A + B • If neither a forward nor reverse reaction is dominant, chemical equilibrium is reached

43. 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

44. Chemistry Comes Alive: Biochemistry Part B 2

45. Biochemistry • Inorganic compounds • Do not contain carbon • Water, salts, and many acids and bases • Organic compounds • Contain carbon, are covalently bonded, and are often large

46. Inorganic: Water • High heat capacity – absorbs and releases large amounts of heat before changing temperature • High heat of vaporization – changing from a liquid to a gas requires large amounts of heat • Polar solvent properties – dissolves ionic substances, forms hydration layers around large charged molecules, and serves as the body’s major transport medium

47. Inorganic: Water • Reactivity – is an important part of hydrolysis and dehydration synthesis reactions • Cushioning – resilient cushion around certain body organs

48. Inorganic: Salts • Inorganic compounds • Contain cations other than H+ and anions other than OH– • Are electrolytes; they conduct electrical currents

49. Inorganic: Acids and Bases • Acids release H+ and are therefore proton donors HCl  H+ + Cl – • Bases release OH– and are proton acceptors NaOH  Na+ + OH–

50. Inorganic: Acid-Base Concentration (pH) • Acidic solutions have higher H+concentration and therefore a lower pH • Alkaline solutions have lower H+ concentration and therefore a higher pH • Neutral solutions have equal H+ and OH– concentrations