Chapter 2 Chemistry

1. Chapter 2Chemistry

2. Unit 4Lecture 1 • Topic: Introduction to Chemistry • Covers Chapter 2 (pg 30 – 32)

3. Recap:Living vs. Nonliving • Differences • Living Organisms: • Made up of at least one cell * Has a metabolism • Has DNA  * Maintains homeostasis • Needs a food source  * Responds to stimuli • Grows * Reproduce • Similarities: • All things (living and nonliving) are made up of MATTER • MATTER - anything that takes up space and has mass • MASS - how much matter an object has

4. Atoms • Matter is made up of chemical elements, or ATOMS • ATOM- basic building block of matter • Smallest, stable unit of matter • An element is a specific type of atom • Elements/atoms cannot be broken down into a simpler stable type of matter • All known elements are arranged into a table • PERIODIC TABLE OF ELEMENTS • Over 100 known elements on Periodic Table, but only around 30 are important to living organisms • 4 Major elements in living organisms: • ~ Oxygen  ~Hydrogen ~ Carbon  ~Nitrogen

7. Atoms • Atoms (elements) are made up of three basic parts: 1. PROTONS  2. NEUTRONS  3. ELECTRONS

8. Atoms • Atomic Mass = Number of Protons + Number of Neutrons • Atomic charge = Number of Protons + Number of electrons • Neutral atoms (atom without a + or – charge) have the same number of electrons as protons • Protons and Neutrons are located in the center of the atom, known as the NUCLEUS • Electrons circle around the nucleus in orbitals • 1st level can hold 2 electrons 2nd level can hold 8 electrons • Element stable when outer orbital(energy level) is full • The only elements that have a full outer orbital are found in the last column of the periodic table • These elements are known as inert gas or noble gas

9. End of Lecture 1

10. Unit 4Lecture 2 • Topic: Types of Bonds • Covers Chapter 2 (pg 33 – 34)

11. Types of Bonds • Most elements are not stable as an individual atom • Elements that are unstable (do not have a full outer orbital) will CHEMICALLY combine with other elements to form a molecule. • When elements chemically combine, it is called a bond

12. Types of Bonds • Types of Bonds: • Covalent Bond – two atoms sharing electrons • Very strong in a watery solution • Example: Carbon dioxide, Oxygen Gas, Water

13. Types of Bonds • Types of Bonds: • Ionic Bond – bond between a positively charged ion and a negatively charged ion (opposites attract) • Ion – an atom with a positive or negative charge • Ionic bond easy to break in a watery solution • When the ionic bond breaks, will go back to a positively charged ion and a negatively charged ion

15. Types of Bonds • Hydrogen Bond • Type of ionic bond that forms between two different water molecules • Very weak, broken easily • But, Hydrogen bonds are very important to living organisms • Causes Cohesion and Adhesion to occur

16. Types of Bonds • Hydrogen Bond • Cohesion – attractive forces between water molecules • EXAMPLE: Surface Tension, Rain Drops • Adhesion – attractive forces between water molecules and another compound/surface • Allows water to move up through narrow tubes against gravity (clings to sides of tubes) • EXAMPLE:Helps plants transport water from roots to leaves

17. End of Lecture 2

18. Unit 4Lecture 3 • Topic: Water, pH, Chemical Reactions • Covers Chapter 2 (pg 35 – 42)

19. Solutions • Solution – mixture of 2 or more substances • Solute – substance dissolved • Solvent – the material dissolving the solute • Aqueous Solution – solution in which water is the solvent • Concentration – measurement of the amount of solute dissolved in the solvent

20. Solutions – Water • Water (H2O) is the universal solvent • Water is formed by covalent bonds between Hydrogen and Oxygen • Polar Molecule – electrons not shared evenly, resulting in a molecule with one side having a negative charge & the other side having a positive charge • The negative charge and positive charge cancel each other out, so the molecule (as a whole) is considered neutral (no charge) • In water, Oxygen has a stronger pull on the electrons • This makes Oxygen slightly negative and the Hydrogensslightly positive  

21. pH Scale • Measuring the concentration of Hydronium Ions (H+) and Hydroxide Ions (OH-) • Scale from 0 – 14

22. pH Scale • Neutral solution • pH = 7; OH = H • Example: water (7), cells (6.5 - 7.5) • Acid • pH < 7; OH– < H+ • More H+ (hydronium) ions than OH– (hydroxide) ions • Sour taste, Highly corrosive • Example: vinegar (3), stomach acid (2), acid rain (<5.6) • Base (aka "Alkaline”) • pH > 7; OH– > H+ • More OH– (hydroxide) ions than H+ (hydronium) ions • Bitter taste, Slippery feel, Soap • Exs: Milk of Magnesia (10.5), Ammonia (11.5), Soap

23. pH Scale • Units on the pH scale are logarithmic • Increase or decrease by factors of 10 • Example: pH 3 is not two times more acidic than a pH 6, but 1,000 times more acidic! • Buffers • Neutralize small amounts of an acid or base • Buffering systems help keep our body’s fluids stay at a normal and safe pH level

24. Energy • Energy– ability to do work or cause change • Comes in many forms, and can change forms • Some types of energy: • Potential, Kinetic, Chemical, Thermal, Solar, Nuclear • Free Energy– energy in a system that is available for work (to fuel cell processes) • Activation Energy – energy required to start a chemical reaction • Catalyst – chemical that reduces amount of activation energy • Enzymes are a main type of catalyst

25. Activation Energy

26. Energy • Na + Cl NaCl •  Reactants   Product(s) • Arrow always points to products • Exergonic Reaction – releases free energy • Endergonic Reaction – absorbs free energy

27. Energy • Life processes require a constant supply of energy • Most common type of cell energy is ATP (Adenosine Triphosphate) • Made up of a 5-carbon sugar, adenine molecule, and a chain of THREE phosphate groups • The phosphate molecules are held together by covalent bonds • When the last phosphate's bond is broken, a lot of energy is released. The energy is used to fuel cell reactions • Forms ADP (Adenosine Diphosphate)

28. End of Lecture 3