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Chapter 2 Chemistry.

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Presentation on theme: "Chapter 2 Chemistry."— Presentation transcript:

1 Chapter 2 Chemistry

2 Unit 4 Lecture 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 Atoms make up everything on Earth!



7 Atoms Atoms (elements) are made up of three basic parts:
1. PROTONS  NEUTRONS  ELECTRONS Protons Neutrons Electrons Charge Positive Neutral Negative Mass Large Very Small

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 Orbitals are known as energy levels

9 End of Lecture 1

10 Unit 4 Lecture 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 4 Lecture 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 Hydrogens slightly 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 ( ) 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

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