Concepts of Matter and Energy

1. Concepts of Matter and Energy By: Kirstin, James, Ryan David, and Chloe

2. Matter • Matter- is anything that occupies space and has weight • Matter can be found in 3 forms • Solid, liquid, and gas • Solids- have a definitive shape and volume • Liquids- have a definitive volume but they conform to the shape of their container • Gases- have neither a definitive shape nor volume • Matter can have both physical and chemical changes • A physical change does not alter the basic nature of the substance • A chemical change do alter the composition of the substance

3. There are 2 types of energy • Kinetic- energy doing work • Potential- inactive stored energy • All living things are built of matter to live and function

4. Forms of energy • Chemical energy- stored in bonds of chemical substance • Electrical energy- movement of charged particles • Mechanical energy-is directly involved in moving matter • Radiant energy-travels in waves, the electromagnetic spectrum

5. Energy forms conversions • Energy can go from one form to another quickly. • Energy can affect body temp. And also the bodies functioning. • When your body heats up is can be a form of Kinetic energy • This chemical reactions in your body is what keeps you alive and functioningeveryday.

6. Composition of Matter Mackenzie, Rolanda, Sam, and Jared

7. Elements and Atoms • Matter is composed of a limited number of substances called elements. • Cannot be broken down, • Total 112 elements, 92 naturally and the rest artificially. • Periodic table is where the complete listing of elements appear. • The atom is the building block of elements and all have atomic symbols on them to show the differences.

8. Atomic Structure • Atoms have unique properties • Protons, electrons, and neutrons, and electrical charge • Protons =positive charge • Electrons= negative charge • Nuetrons=neutral charge • Electrical charger= able to attract and repel all the different properties. • Which all this together in one atom weighs at 1 atomic mass unit.

9. Planetary and orbital models of an Atom • Planetary model portrays the atom as a miniature solar system. • Orbital model the more modern model of the atomic structure. • The atomic nucleus includes the protons and neutrons • The nucleus is dense and positively charged • The elements orbit the nucleus • The orbital's are the regions around the nucleus where you would find the electrons.

10. Identifying elements • Protons, neutrons, and electrons are all a like. • Atoms of different elements= have different electrons, protons, and neutrons • All known atoms can be described by adding one proton and one electron. • Light protons have equal number of protons and neutrons larger have more neutrons then protons.

11. Atomic number • All elements have an atomic number • Atomic number is equal to the number of protons its atoms contain. • Each atom has different number of protons than atoms of any other element. • Protons are always equaled to the number of electrons. • The atomic number indirectly tells the number of electrons the atoms has.

12. Atomic Mass Number • Atomic mass number of any atom is the sum of the protons and neutrons contained in its nucleus • ( EX. Hydrogen has one bare proton and no nuetrons. So its atomic number and atomic mass number is the same which is 1)

13. Atomic Weight and Isotopes • Atomic weight is the average of the mass numbers of all the isotopes of an element. • Isotopes have the same number of protons and electrons but vary in the number of neutrons they contain. • Radioistopes are havire isotopes of atoms that are unstable and tend to decompose to become more stble. • Radioactivity is the process of spontaneous atomic decay. It can be compared to a tiny exposions • Isotopes of an element have the same atomic number but different atomic masses. • The atomic weight of an element is approximately equal to the mass number of its abundant isotope

14. Michael MoussalliPatrick McGinnSydney O’Connor Molecules and Compounds Pages 32-33

15. Definitions • Molecule – When 2 or more of the same atoms combine chemically • H + H H2 (Molecule) • Compound – When 2 or more different atoms bind together to form a molecule • 4H + C CH4 (Compound)

16. Molecular Formula • The formula that shows what different atoms make a molecule or compound • H + H • 2H + O

17. Chemical Equation • The whole equation • H + H = H2 • 2H = O = H2O

18. Chemical Bonds and Chemical Reactions Abby, Kristina, Nick, and Kenneth

19. What is a Chemical Reaction? A chemical reaction occurs when atoms combine or separate from other atoms.

20. Bond Formation Chemical bonds are not physical structures. They are energy relationships that involve interactions between electrons of reacting atoms.

21. Roles of Electrons Electron shells- (energy levels) fixed regions of space around the nucleus where electrons are found. These shells are numbered 1 to 7, starting from the nucleus outward. The currently known maximum number of electron shells is seven. The closest to the nucleus are the most strongly positive charged. The more distant are more likely to react with other atoms(less securely held).

22. Roles of Electrons (Cont.) Each shell has a limit to how many electrons it can hold. (Shell 1 can hold 2, shell 2 can hold 8, etc.) Valence Shells- the outermost shell in the atom. It determines the chemical behavior of the atom. The electrons in inner shells do not participate in bonding.

23. Ionic Bonds Ionic bonds form when electrons are completely transferred from one atom to another. They gain or lose electrons during bonding. This causes their positive and negative charges to become unbalanced which result in ions, charged particles.

24. Ionic Bonds (Cont.) Negatively charged ions are called anions. Positively charged ions are called cations. Both result when an ionic bond is formed. Most bonds formed under ionic bonds fall under the category of salts.

25. Covalent Bonds • Covalent bonds occur when covalent molecules share electrons. • In some covalent molecules the electrons are shared equally between the atoms of a molecule. These are called Nonpolar Covalently Bonded Molecules. • When covalent bonds are made, the molecule formed always has a definite three-dimensional shape. (PG 35) • The shape determines what other molecules or atoms in can interact with. • The shape may also result in unequal electron-pair sharing.

26. Covalent Bonds (Cont.)

27. Hydrogen Bonds Hydrogen bonds are extremely weak bonds formed when a hydrogen atom bound to one electron-hungry nitrogen or oxygen atom is attracted by another electron-hungry atom. The hydrogen atoms form a bridge between the atoms. Hydrogen bonds are common in water molecules (surface tension). Hydrogen bonds also can be called Intramolecular bonds. These bonds help to bond different parts of the same molecule together into a special three-dimensional shape.

28. Patterns of Chemical Reactions • Chemical reactions involve the formation or destruction of bonds between atoms. • There are three types of chemical reactions: • Synthesis • Decomposition • Exchange

29. Synthesis Reaction • A synthesis reaction occurs when 2 or more atoms combine to form a larger more complex molecule. A + B AB • Synthesis reactions always involve bond formation. • They are energy-absorbing reactions. • Synthesis reactions underlie all anabolic activities that occur in body cells. • They are important for growth and repair of old tissue.

30. Decomposition Reaction • A decomposition reaction occurs when a molecule is broken down into smaller molecules, atoms, or ions. AB A + B • Decomposition underlie all catabolic processes in the body cells. • Decomposition Reactions are basically Synthesis Reactions in reverse. • Bonds are always broken and the products of the reaction are smaller. • Chemical energy is released.

31. Exchange Reaction • Exchange reactions involve both synthesis and decomposition reactions, both bonds are made and broken. • A switch is made during the reaction between molecule parts. AB + C AC + B and AB + CD AD + CB

32. Biochemistry: The Chemical Composition of Living Matter Ann Marie Dean, Brett Foxworth, Casey Gurlaskie, Chase Williams

33. 2 major classes of molecules in the body: • Inorganic and Organic • Inorganic compounds: Lack Carbon; tend to be small and simple-ex: water, salts • Organic compounds: Carbon-containing compounds; more complex-ex: carbohydrates, lipids, proteins, nucleic acids

34. Inorganic Compounds • Water • Most abundant in the body (2/3) • Properties that make water so vital: • High heat capacity- absorbs and releases large amounts of heat before its temperature changes appreciably • Polarity/solvent properties-water=universal solvent • Solvent=liquid or gas in which small amounts of other substances (solute) can be dissolved • Solvent+solute=solution • Respiratory gases and waste can dissolve in water

35. A. Water • Mucus and saliva use water as their solvent 3. Chemical Reactivity- • Water=very important type of reaction • Water molecules are added to larger bonds to digest food and break down biological molecules (hydrolysis reaction) 4. Cushioning-water forms a cushion around organs, such as the brain, to protect it from trauma

36. B. Salts • Most plentiful salts contain calcium and phosphorus (found in bones and teeth) • Dissociation: when salt dissolves in the body and separates in its ions • Salt is vital to a body functioning • All salts are electrolytes-substances that conduct and electric current in a solution

37. C. Acids and Bases • Characteristics of Acids: • Sour taste, can dissolve metal • Acid-a substance that can release hydrogen ions in detectable amounts • Acids are proton donors • Acids in the body: hydrochloric, acetic, carbonic • Strong vs. weak acids

38. C. Acids and Bases 2. Characteristics of Bases: • Bitter taste, slippery • Proton accepters • Hydroxides are common inorganic bases • Any base containing the hydroxl ion is a strong base • Bicarbonate ion (in blood) is weaker • When acids and bases mix, they form water and salt • Neutralization reaction-when an acid and a base react

39. C. Acids and Bases 3. PH: Acid-Base concentrations • The relative concentration of hydrogen ions in various body fluids is measured in concentration units called PH units • PH scale: found in 1990 by Sorensen and is based on the the # of protons in a solution expressed in terms of moles per liter • PH scale-014; 7= midpoint (neutral); 1-6 acids, 8-14 bases • 1= stronger acid; 14=weaker base • Living cells are extraordinarily sensitive to even slight changes in pH • Acid/base balance is regulated by the kidney, lungs, and a # of chemicals called buffers (present in body fluids)

40. C. Acids and Bases • Weak acids and weak bases are important to the body’s buffers systems which act to maintain pH stability • Regulation of blood pH is critical • Normal blood pH ranges from 7.35-7.45; small changes in this threaten death • When blood pH begins to dip into the acid range, the amount of life-sustaining oxygen that the hemoglobin in blood can carry to body cells begins to fall rapidly to dangerously low levels

41. Acids and Bases

42. pH Scale