1. Chemical Reactions Balancing Equations

2. How do you know a chemical reaction has occurred? A change in matter that released energy as both heat and light is strong evidence that a chemical reaction has occurred. The evolution of gas when two substances are mixed is often evidence of a chemical reaction.

3. How do you know a chemical reaction has occurred? A solid that is produced as a result of a chemical reaction in solution that separates from the solution is known as a precipitate. A change in color is often an indication of a chemical change.

4. Chemical Equations An equation that represents the identities and relative amounts of the reactants and products in a chemical reaction. All products and reactants must be identified. The law of conservation of mass must be obeyed!

5. Rules for Balancing Equations 1.Balance the different types of atoms at one time. 2.First balance the atoms of elements that are combined and only appear once on each side of the equation. 3.Balance polyatomic ions that appear on both sides of the equation as a single unit. 4.Balance H atoms and O atoms after atoms of all other elements have been balanced.

6. Types of Reactions: Synthesis or Combination A + X  AX A + X  AX A chemical change in which two or more substances react to form a single new substance. 2Mg (s) + O 2(g)  2MgO (s) 2Mg (s) + O 2(g)  2MgO (s) 2Na (s) + F 2(g)  2NaF (s) 2Na (s) + F 2(g)  2NaF (s) SO2 (s) + H 2 O (l)  H 2 SO 3(aq) SO2 (s) + H 2 O (l)  H 2 SO 3(aq)

7. Types of Reactions: Decomposition AX  A + X A chemical change in which a single compound breaks down into two or more simpler products. 2HgO (s)  2Hg (l) + O 2(g) 2HgO (s)  2Hg (l) + O 2(g) Ca(OH) 2(s)  CaO (s) + H 2 O (g) Ca(OH) 2(s)  CaO (s) + H 2 O (g)

8. Types of Reactions: Single-Replacement A + BX  AX + B Y + BX  BY + X A chemical change in which one element replaces a second element in a compound. 2Na (s) + 2H 2 O (l)  2NaOH (aq) + H 2(g) 2Na (s) + 2H 2 O (l)  2NaOH (aq) + H 2(g) Cl 2(g) + 2KBr (aq)  2KCl (aq) + Br 2(l) Cl 2(g) + 2KBr (aq)  2KCl (aq) + Br 2(l) F 2(g) + 2NaCl (aq)  2NaF (aq) + Cl 2(g) F 2(g) + 2NaCl (aq)  2NaF (aq) + Cl 2(g)

9. Types of Reactions: Double-Replacement AX + BY  AY + BX A chemical change involving an exchange of positive ions between two compounds. 2KI (aq) + Pb(NO 3 ) 2(aq)  PbI 2(s) + 2KNO 3(aq) 2KI (aq) + Pb(NO 3 ) 2(aq)  PbI 2(s) + 2KNO 3(aq) HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l) HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)

10. Types of Reactions: Combustion Hydrocarbon + O 2  CO 2 + H 2 O A chemical change in which an element or compound reacts with oxygen often producing energy in the form of heat and light. C 3 H 8(g) + 5O 2(g)  3CO 2(g) + 4H 2 O (g) C 3 H 8(g) + 5O 2(g)  3CO 2(g) + 4H 2 O (g)

11. Liquids vs. Solids Physical propertyLiquidsSolids CompressibilityAlmost incompressible, but expand slightly when heated Almost incompressible due to tightly packed particle arrangement DensityLiquid particles are free flowing, causing them to be less dense than solids Due to the closely packed particle arrangement, solids have a high density ShapeTake the shape of the container Rigid shape Rate of diffusionParticles are free flowing, so they diffuse faster than a solid but slower than a gas Do not diffuse

12. Three States of Matter

13. Intermolecular Forces  Dispersion forces are very weak forces of attraction between molecules resulting from momentary dipoles occurring due to uneven electron distributions in neighboring molecules as they approach one another.  Dipole-dipole Interactions are stronger intermolecular forces than Dispersion forces that occur between molecules that have permanent net dipoles (polar molecules). polar  Hydrogen bonds occur between molecules that have a permanent net dipole resulting from hydrogen being covalently bonded to either fluorine, oxygen or nitrogen. dipole

14. Intermolecular Forces

15. Chemical Bonds Covalent bonds result from electrons being shared between atoms in a molecule. Polar- A bond between atoms in which the electrons are unequally shared. Polar- A bond between atoms in which the electrons are unequally shared. Non-polar-A bond in which electrons are equally shared by two atoms. Non-polar-A bond in which electrons are equally shared by two atoms.

16. Viscosity Viscosity is the friction between moving molecules in a liquid. The stronger the intermolecular forces, the higher the viscosity. The stronger the intermolecular forces, the higher the viscosity. Because hydrogen bonding is the strongest intermolecular force, liquids with hydrogen bonding result in very high viscosities. Because hydrogen bonding is the strongest intermolecular force, liquids with hydrogen bonding result in very high viscosities.

17. Surface Tension An inward force that tends to minimize the surface of a liquid causing the surface to behave as a thin skin. Water has a higher surface tension than most liquids due to its strong hydrogen bonding. Water has a higher surface tension than most liquids due to its strong hydrogen bonding. This property allows insects to walk on water. This property allows insects to walk on water.

18. Capillary Action Capillary action is the tendency of a liquid to rise in narrow tubes or to be drawn into small openings. Water is able to rise in plants’ roots and stems because of this property. Water is able to rise in plants’ roots and stems because of this property.

19. Vapor Pressure Vapor pressure is the result of molecules escaping from the surface of the liquid and entering the vapor phase. Because hydrogen bonds hold water molecules to each other, the tendency of these molecules to escape is low. Because hydrogen bonds hold water molecules to each other, the tendency of these molecules to escape is low. Water has an unusually low vapor pressure. Water has an unusually low vapor pressure.

20. Properties of Solids  The smallest repetitive unit in a crystalline structure is called a unit cell.  Some solids, such as amorphous solids lack an ordered internal structure. These solids may morph over a wide range of temperatures before melting These solids may morph over a wide range of temperatures before melting  In a molecular solid, intramolecular covalent bonds are stronger than intermolecular attractive forces.

21. Solids

22. Solutions

23. Solutions homogeneous mixtures can be atoms, ions or formula units Do not separate upon standing Do not scatter light Suspension: particles remain thoroughly mixed while the liquid is being stirred particles will settle to the bottom over time w/o agitation Can be separated by filtration ColloidsHeterogeneous Do not separate upon standing or filtration Scatter light (Tyndall effect)

24. Solutions What affects whether a substance will dissolve? nature of the solvent (substance dissolving the solute) nature of the solute (substance dissolved in the solvent) What factors determine how fast a soluble substance dissolves?

25. Solutions Agitation (stirring): makes the solute dissolve faster does not affect the amount of solute that dissolves does not affect the amount of solute that dissolves Temperature: Temperature: dissolves quicker in hot water than cold water kinetic energy of the water molecules is greater at the higher temperature Particle size: powder dissolves quicker than a single crystal due to greater surface area of ions being exposed to the colliding water molecules powder dissolves quicker than a single crystal due to greater surface area of ions being exposed to the colliding water molecules

26. Solutions Miscible: two liquids or gases that will dissolve in each other can mix in any proportion example: water and ethanol example: water and ethanolImmiscible: liquids or gases that are insoluble in one another Example: oil and water

27. Solutions Solubility The maximum amount of a substance that dissolve in a given quantity of solvent at a constant temperature and pressure to produce a saturated solution. Is NaCl soluble in water? Does it matter how much NaCl that you add to a specific amount of water?

28. Solution saturated solution: contains the maximum amount of solute for a given amount of solvent at a given temperature contains the maximum amount of solute for a given amount of solvent at a given temperature unsaturated solution: contains less solute than specified by the solubility at a given temperature supersaturated: contains more than the standard amount of solute specified by the solubility at a given temperature

29. Solutions Solubility increases with temperature generally Exception: solubility of gases decreases with increase in temperature The solubility of a gas increases as the partial pressure of the gas above the solution is increased Ex. Carbonated beverages