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Chemistry List the first five things that come to mind when you hear this word.

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Presentation on theme: "Chemistry List the first five things that come to mind when you hear this word."— Presentation transcript:

1 Chemistry List the first five things that come to mind when you hear this word.

2 Vocabulary: Define as many of the following words as you can: Atomic Mass Atomic Number Electron Ion Neutron Proton Isotope Valence (Electrons) Covalent Bond Ionic Bond Molecule Isomer

3 The Periodic Table

4 Three Subatomic Particles Proton: (+) charged particle found inside the nucleus Neutron: neutral particle found inside the nucleus Electron: (-) charged particle found outside the nucleus in various energy levels

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6 Atomic Number & Atomic Mass Atomic Number: the number of protons (and electrons in neutral atoms) Mass Number: the number of protons plus the number of neutrons (mass of nucleus) Atomic Mass (Weight): weighted average of the isotopes for an element Mass # Atomic #

7 Isotopes of Hydrogen What are isotopes? Atoms of the same element with the same # of protons, different # of neutrons

8 The Periodic Table Groups/Families (columns) Periods (rows) Atomic Number Atomic Mass Valence Electrons

9 Valence Electrons: electrons in outermost energy level Most important for bonding

10 Covalent Bonds Valence electrons are shared between atoms

11 Covalent Bonding Sharing one or more electron pairs between 2 atoms

12 Covalent Bonds

13 Ions Atoms of the same element with a different number of electrons (charged particles) Cations (positive ions) – lose electrons Anions (negative ions) – gain electrons

14 CATION “cat”ion ca+ion ANION “ant”ion n - negative

15 Ionic Bonds Valence electrons are transferred between atoms

16 Ionic bonding

17 Ionic Bonding

18 Molecule Smallest unit of most compounds that displays all the properties of that compound

19 Counting Atoms Na 2 SO Na, 1 - S, 4 - O 7 atoms Ca(OH) Ca, 2 - O, 2 - H 5 atoms 3 Fe 2 (SO 3 ) Fe, 9 - S, 27 - O 42 atoms

20 Isomers Compounds with the same molecular formula but different structures (arrangement of atoms) and, therefore, different properties

21 Section 2.2

22 What do you know about water?

23 Unique Properties of Water: Most abundant compound in nearly all living things

24 Unique Properties of Water: Density of solid water vs. liquid water Expansion on freezing – water forms a crystalline structure that expands and is less dense than its liquid state

25 Unique Properties of Water: Covers more than 75% of the Earth’s surface Water 100°C or 212°F Water 0°C or 32°F Universal solvent

26 Water has a high heat capacity Water has a high heat capacity Heat capacity = the amount of heat energy required to increase its temperature, is relatively high. Large bodies of water (oceans and lakes) can absorb large amounts of heat with only small changes in temperature. This protects organisms living within from drastic changes in temperature. Unique Properties of Water:

27 What is a polar molecule? An uneven distribution of electrons The molecule becomes charged on each end It means that the electrons are NOT shared equally

28 Hydrogen Bonding Because of their partial positive and negative charges, polar molecules such as water can attract each other. The attraction between a hydrogen atom on one water molecule and the oxygen atom on another is known as a hydrogen bond.

29 Interaction Between Water Molecules Negative Oxygen atomof one water molecule is attracted to the Positive Hydrogen atom of another water molecule to form a HYDROGEN BOND

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31 Cohesion Cohesion is when one water molecule sticks to another water molecule  due to hydrogen bonding It creates surface tension (a measure of strength) at the surface of the liquid Produces a surface “film” that allows small insects to “walk on water”

32 Surface Tension A measure of how difficult it is to break the surface of a liquid because of hydrogen bonds. The surface tension of a liquid results from an imbalance of intermolecular attractive forces, (cohesive forces) between molecules A molecule in the bulk liquid experiences cohesive forces with other molecules in ALL directions. A molecule at the surface of a liquid experiences only NET INWARD cohesive forces.

33 Adhesion Adhesion is when a water molecule sticks to another substance (NOT another water molecule) A piece of tape adheres to the wall Capillary action is an example of adhesion (water moving from the roots, through the stem, & into the leaves of a flower)

34 Capillary Action Water climbing up a small tube because of the adhesion to the sides and the cohesion between water molecules.

35 Adhesion Also Causes Water to … hold onto plant leaves Attach to a silken spider web

36 Cohesion and Adhesion When liquid water is confined in a tube, its surface (meniscus) has a concave shape because water adheres to the surface and creeps up the side. Mercury does not wet glass - the cohesive forces within the drops are stronger than the adhesive forces between the drops and glass. When liquid mercury is confined in a tube, its surface (meniscus) has a convex shape because the cohesive forces in liquid mercury tend to draw it into a drop.

37 High Heat Capacity Water’s heat capacity, the amount of heat energy required to increase its temperature, is relatively high. Because of the multiple hydrogen bonds between water molecules, it takes a large amount of heat energy to cause those molecules to move faster and raise the temperature of the water. Large bodies of water, such as oceans and lakes, can absorb large amounts of heat with only small changes in temperature. This protects organisms living within from drastic changes in temperature.

38 Universal Solvent Hydrophobic A substance that repels (is afraid of) water (a nonpolar molecule) Examples: oil Hydrophilic A substance that is attracted to water (a polar molecule) Examples: sugar, alcohol The Chemistry of Water

39 Solutions and Suspensions

40 What is a Mixture? Substance composed of two or more elements or compounds that are mixed together, but NOT chemically combined. Mixture Examples: salt & pepper; sugar & sand; salt & sand; Earth’s atmosphere; soil

41 Solutions and Suspensions Mixture – a material that consists of 2 or more elements or compounds physically mixed together but are not chemically combined.

42 Homogeneous Mixture (Solution) Homogeneous mixture in which one substance is dissolved in another Examples: sugar in water; salt in water; hot chocolate; tea, Kool-aid Salt Dissolving in Water

43 Parts of a Solution All solutions contain a solute and a solvent. Solvent is the substance doing the dissolving. Example: water Solute is the substance that is being dissolved. Example: salt

44 Heterogeneous Mixture (Suspensions) Mixture of water and non-dissolved materials (heterogeneous) Examples: blood; oil and water

45 Application: 1 – Choose one of your examples of a solution. Explain how the parts (components of the solution) can be separated. 2 – For the same example, identify the solute and solvent in the solution. Explain your reasoning.

46 Example of a Solution: The solute (sugar) dissolves in the solvent (water). The solute (Kool-aid) dissolves in the solvent (water). The solute (oxygen) dissolves in the solvent (nitrogen). Water is the universal solvent.

47 Homeostasis Ability to maintain a steady state despite changing conditions Water is important to this process because… a. Makes a good insulator b. Resists temperature change c. Universal solvent d. Coolant e. Ice protects against temperature extremes (insulates frozen lakes)

48 All about the pH Scale

49 The Water Molecule: Water dissociation animation

50 Acids Compounds that release H + into solution H + = hydrogen ion If [H+] > [OH-] then the solution is acidic. HCl  H + + Cl -

51 Bases Compounds that release OH - into solution OH - = hydroxide ion If [OH-] > [H+] then the solution is basic NaOH  Na + + OH -

52 The pH scale (power of Hydrogen) as in Hydrogen ions Measures the relative concentration of H + and OH -

53 Interpreting the pH Scale pH is based on a logarithmic scale A pH of 5 is 10× more concentrated than a pH of 6 (in terms of H + concentration) A pH of 9 is 100× more concentrated than a pH of 11 (in terms of H + concentration)

54 Classifying Acids & Bases Acids: Strong acids have a pH value between 1- 3 Weak acids have a pH value between 4- 6 Bases: Strong bases have a pH value between Weak bases have a pH value between 8- 11

55 Buffers Weak acids or bases that can react with strong acids or bases to prevent sharp sudden changes in pH Produced naturally by the body to maintain homeostasis The pH value in most cells is The pH of stomach acid is 2 The pH of the blood is 7.4

56 Blood Buffer – pH needs to be 7.4 Carbonic Acid (Lungs – carried away as CO2) bicarbonate ion (buffer) (Kidney – carried away in urine) Double arrows mean reactions are in equilibrium Free hydrogen ions

57 Blood Buffer – pH needs to be 7.4 You can see that if the reactions go to the right, hydrogen ions are released, making the solution more acidic, and if the reactions go to the left, hydrogen ions are sucked up, making the solution more basic. The question that is probably eating away at you right now is: "how do I know which way the reactions are going to go-- to the right or to the left?“ Good question! The bicarbonate ion isn't a very strong acid or base. It doesn't have to go one way or the other all the time. Instead, the direction it goes depends on the solution it is in. You see, if it is in an acidic solution, there are lots of hydrogen ions floating around; in this situation, the presence of tons of hydrogen ions will force the reaction to go to the left. Do you see why? Look at the reaction again. All you need for the reaction to be able to go to the left is available hydrogen ions (see the H+ in the middle that is needed for the reaction to go to the left?). So if bicarbonate ions find themselves in acidic solutions, they tend to act like the weak base they can be, and suck up the excess hydrogen ions. Consider what will happen if bicarbonate is put into a basic solution. A basic solution has very few hydrogen ions floating around. In this condition, the bicarbonate reaction cannot proceed to the left, since no hydrogen ions are available. Instead, it will go to the right, producing hydrogen ions. In this manner, bicarbonate ions will tend to act like the weak acid they can be, and release hydrogen ions.

58 1. Use the structure of a water molecule to explain why it is polar. 2. Why is water such a good solvent? 3. What is the difference between a solution and a suspension? 4. How do hydrogen bonds between water molecules occur? 5. How does the body counteract a drop in blood pH during strenuous exercise? 6. If a strong acid is dissolved in pure water will the pH of the solution be greater or less than 7?

59 Water Chemistry 1. Please draw the structure of a water molecule (H 2 O). 2. Label the charges on the molecule that make it polar (negative and positive ends). 3. See if you can connect the water molecule that you drew to two other molecules of water using hydrogen bonds. 4. Can you think of any unique properties that water displays?

60 Water Chemistry 1. Explain why water is “polar” in your own words. 2. Explain capillary action in your own words. 3. Explain surface tension in your own words.

61 1. According to the scale, what substance is neutral? 2. What solution has an [H+] of 0.01? 3. Which substance produces the most OH- ions in solution? 4. How many more times basic is pH 14 compared to pH 11.

62 Chemical Reactions and Enzymes Section 2.4

63 Objectives  Distinguish between potential and kinetic energy.  Differentiate between reactants and products in a chemical reaction.  Explain the role of a catalyst in chemical reactions.

64 What is energy?  The ability to do work or cause a change.  Can exist in different forms: thermalradiant chemical nuclear mechanical electrical

65 Forms of Energy:  POTENTIAL ENERGY  stored energy  KINETIC ENERGY  energy of motion  FREE ENERGY  the energy available to do work

66 Activation Energy  the energy needed to start a chemical reaction.

67 Chemical Reactions:  Everything that happens in an organism is based on chemical reactions.  A chemical reaction is a process that changes one set of chemicals into another set of chemicals. CO 2 + H 2 O  C 6 H 12 O 6 + O 2

68 Chemical Reactions:  The elements or compounds that enter into the reaction are the reactants.  The elements or compounds produced by the reaction are the products  Bonds are broken/formed in chemical reactions.

69 What are the Reactants? Products? CO 2 + H 2 O  C 6 H 12 O 6 + O 2 ReactantsProducts  means “Yields” or “Produces”

70 Energy RELEASED Reactions  Activation Energy : the energy needed to start a chemical reaction.  Exergonic (Exothermic) Reactions - involve a net release of free energy  Often happen spontaneously

71 Energy ABSORBED Reactions  Activation Energy : the energy needed to start a chemical reaction.  Endergonic (Endothermic) Reactions - involve a net absorption of free energy  Will not occur without a source of energy

72 Enzymes  Enzymes – are proteins that act as a biological catalyst.  Catalyst – a substance that speeds up a chemical reaction by lowering the activation energy. Draw the enzyme pathway on your graph

73 Enzyme Action Substrate – the reactants of an enzyme catalyzed reaction Active Site – where the substrate and enzyme join.  Have complimentary shapes  Can only bind with a specific molecule – “induced fit”

74 How Enzymes Work:  Enzymes act on a specific substrate. (Shapes fit together like a lock and a key)  A small area on the enzyme, called the active site, can attract and hold only a specific substrate.  The enzyme acts as a biological catalyst, which accelerates the rate of the chemical reaction.  The enzyme reduces the activation energy needed by weakening the chemical bonds in the substrate.  The enzyme is then released unchanged.

75 Effects of Enzymes:

76 Regulation of Enzyme Activity  Enzymes work best at certain pH levels and temperatures.  Most enzymes in humans work best at 37  C  Denaturation – a change in the enzymes shape  Enzyme becomes nonfunctional  Heat and pH can cause denaturation  Some enzymes can be turned on and off  Lactose intolerance

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78 1. What letter represents the product? 2. What letter represents the reactant? 3. What letter represents the activation energy? 4. How should the x and y axis be labeled? 5. Is the reaction exergonic or endergonic? Explain


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