1. Chapter 2

2.  Introduction to “Baby Chemistry”  Life depends on chemistry  We need to be able to understand the chemistry before we move on to understand the rest of body

3.  Atom - the basic unit of matter  Atoms are made of 3 subatomic particles 1. Protons - positive charge 2. Neutrons – neutral charge 3. Electrons – negative charge

4.  Protons and neutrons are found in the atoms nucleus - The center of the atom  Electrons are found outside the nucleus  Electron cloud  They are in constant motion  Are attracted to positive charge of nucleus but remain outside due to energy of motion

5.  What is the overall charge of the nucleus?  What is the overall charge of the electron cloud?  Nucleus (+) & Electron Cloud (-)  Therefore, atoms are neutral from equal numbers of protons and electrons

6.  Elements – pure substances that consist of entirely 1 type of atom  They are represented by 1, 2 or 3 letters  There are 117 elements (periodic table in back of book)  Based on certain characteristics of each element

7.  Periodic table includes information on atomic mass and atomic number  Atomic Number – number of protons (and therefore electrons)  Atomic Mass – sum of protons and neutrons

9.  How are isotopes different?

10. Q) How are isotopes different? A) They have a different number of neutrons  Isotopes are identified by their atomic mass number (protons + neutrons)  All isotopes still have the same chemical properties (because neutrons are neutral)

11.  Have nuclei that are unstable and break down at constant rates over time Important Uses 1. Determine age of rocks 2. Treat cancer and kill bacteria that cause food to spoil 3. Used as “tracers” to follow substances in the body

12. *In nature, most elements are found combined with other elements in compounds Compound – substances formed by the chemical combination of 2 or more elements

13.  What information is contained in a chemical formula? 1. Type of atoms 2. How many of those atoms  Example) Q) What’s in H 2 0? A) 2 hydrogen and 1 oxygen

14.  Physical and chemical properties of a compound are usually different from those of elements from which they are formed  Example – NaCl  Sodium Chloride : Table Salt  Sodium (Na) – silver colored metal  Chlorine (Cl) – greenish gas +=

15.  Hold compounds together  Involves the outer electrons (valence electrons) that surround each atom’s nucleus  3 types of bonds 1. Ionic (strongest) 2. Covalent 3. Van Der Waals Forces (weakest)

16.  One or more electron is “transferred” from one atom to another  An atom that loses an electron has a positive charge  An atom that gains an electron has a negative charge  Positively or negatively charged atoms are called ions

18.  When electrons are “shared” instead of being “transferred”  How do they share electrons?  The electron travels in the orbits of both atoms  Single bond – 2 electrons shared  Double bond – 4 electrons shared  Triple bond – 6 electrons shared

19.  Atoms joined by covalent bonds are called molecules  Water Molecule

20.  Atoms in covalent bond don’t always share equally  Rapid movement of electrons creates “+” and “-” charges  When molecules are close together, an attraction can occur between oppositely charges region  Creates intermolecular “Vander Waals Forces”

21.  How does the gecko’s foot function similarly to Van Der Waals Forces?

22.  The water molecule – H 2 O (think mickey mouse)

23.  Water is a polar molecule  What is polarity?  The uneven distribution of electrons which causes the molecule to have regions with a positive and negative charge  Allows water to attract to other molecules

24.  Most abundant compound in living things  Naturally occurring and found at any temperature on the earth  Water expands when it freezes, SO WHAT??  How does that affect us?

25.  Held together by hydrogen bonds (not as strong as ionic or covalent) which form between the hydrogen and oxygen  Why would this be important?  How many bonds are there per water molecule?

26.  Allows 4 bonds to form  This bonding pattern is what makes water responsible for it’s many special properties

27.  Cohesion  An attraction between water molecules  Why can a spider walk on water  How does water come out of a faucet?

28.  Adhesion  An attraction of water to other molecules  How does a tree get water from its roots to its leaves?  Why does the water level in a graduated cylinder look curved?

29.  Mixture – a material composed of 2 or more elements or compounds that are physically combined, but not chemically combined  Examples:  Chocolate Chip Cookie  Earths Atmosphere  2 types of Mixtures 1. Solution 2. Suspension

30.  Is a mixture of 2 or more substances in which the molecules are evenly distributed  Example:  Salt dissolving in water  (the NaCl and H 2 O combine by the Cl - being attracted to the H and the Na + being attracted to the O)

31.  2 important terms to know  Solute – substance being dissolved  Solvent – substance doing the dissolving Q) Why is water considered the greatest solvent in the world? A) Polarity allows it to pull other molecules apart

32.  Mixture of water and non-dissolved material

33.  Concentrations for hydrogen (H + ) and hydroxide (OH - ) ions are important indicators of the properties of a solution  A measurement system was created to indicate the concentration of the H + ions in the solution called the pH scale  pH – potential of hydrogen  A pH of 7 = equal number of H + and OH - ions

34. H + + OH - = H 2 O  Solutions with a pH < 7 are called acids  Because more hydrogen ions (H + )  The lower the pH, the greater the acidity

35.  Solutions with a pH > 7 are called alkaline or basic  Because more hydroxide (OH - ) ions  The higher the pH, the more basic the solution

36.  Each step of the pH scale represents a factor of 10  Ex) pH of 4 has 10x as many H + ions as a solution with a pH of 5 AcidsBases I-------------------------I-------------------------I 1 H + > OH - 7 H + < OH - 14 H + = OH -

37.  Weak acids or bases that can react with strong acids or bases to prevent sharp, sudden changes in pH  Why are buffers important?  The pH of fluids within most cells of the human body must generally be kept between 6.5 and 7.5  If pH is higher/lower – it will affect the chemical reactions that take place  Controlling the pH of cells is important for maintaining homeostasis