1. IIIIIIIVV Ch. 10 – Chemical Reactions

2. A. Signs of a Chemical Reaction Evolution of heat and light Formation of a gas Formation of a precipitate (solid) Color change

3. B. Chemical Equations Have two parts: Have two parts: 1.Reactants = the substances you start with 2.Products = the substances you end up with The reactants will turn into the products. A + B  C + D A + B  C + D Reactants  Products

4. B. Chemical Equations

5. Symbols used in equations ■Special conditions other than heat can be written over the arrow also. ■Such as “high pressure” ■Or a catalyst – which makes the reaction go faster, but is not changed in the reaction. Hi Pressure MnO 2

6. Example: Translating a reaction into an equation In the catalytic converter of a car, carbon monoxide gas and oxygen gas become carbon dioxide in the presence of Platinum metal ** special condition= catalyst! CO(g) + O 2 (g) CO 2 (g) CO(g) + O 2 (g) CO 2 (g) Pt

7. Law of Conservation of Mass mass is neither created nor destroyed in a chemical reaction 4 H 2 O 4 H 2 O 4 g32 g 36 g total mass stays the same atoms can only rearrange

8. Describing Balanced Equations to produce How many? Of what? In what state? Zn(s) + 2HCl(aq)  ZnCl 2 (aq) + H 2 (g) One atom of solid zinc reacts with two molecules of aqueous hydrochloric acid one unit of aqueous zinc chlorideand one molecule of hydrogen gas.

9. Describing Equations Describing Coefficients: –individual atom = “atom” –covalent substance = “molecule” –ionic substance = “unit” 3 molecules of carbon dioxide 2 atoms of magnesium 4 units of magnesium oxide 3CO 2  2Mg  4MgO 

10. Counting Atoms To determine if an equation is balanced or not, you need to be able to count atoms in a compound. Cu(NO 3 ) 2 has.... 1 copper atom 2 nitrogen atoms 6 oxygen atoms

11. IIIIIIIVV Chapter 10 “Chemical Reactions” Balancing Reactions

12. Balanced Chemical Equations Atoms can’t be created or destroyed in an ordinary reaction: –All the atoms we start with we must end up with (meaning: balanced!) A balanced equation has the same number of each element on both sides of the equation.

13. Examples of balanced chemical Equations 2 Ag + S  Ag 2 S CH 4 + 2 O 2  CO 2 + 2 H 2 O  + + + 

14. How to Balance Chemical Equations You can only balance equations using COEFFICIENTS in front of a reactant or product!!! Never change the SUBSCRIPTS in a reactant or product!!!! –If you change the subscript (formula) you are describing a different chemical. –H 2 O is a different compound than H 2 O 2 Never put a coefficient in the middle of a formula; they must go only in the front 2NaCl is okay, but Na2Cl is not. 2NaCl is okay, but Na2Cl is not.

15. Tips for Balancing Equations Start with elements that only appear once on BOTH sides of the equation CO + O 2  CO 2 Balance Oxygen and Hydrogen last… they usually appear more than once in an equation 2CO + O 2  2CO 2

16. If a polyatomic ion appears unchanged on both sides of the equation, treat it as a single unit, instead of counting individual atoms CuSO 4 + NaCl  Na 2 SO 4 + CuCl 2 CuSO 4 + 2NaCl  Na 2 SO 4 + CuCl 2 Check to see if the coefficients used give equal numbers of atoms on both sides of arrow. If not, go back and re-balance!!! Last, if all the coefficients are multiples of each other, reduce to the least common multiple.

17. Practice Balancing Examples _AgNO 3 + _Cu  _Cu(NO 3 ) 2 + _Ag _Mg + _N 2  _Mg 3 N 2 _P + _O 2  _P 4 O 10 __CaH 2 + __H 2 O  _H 2 + __Ca(OH) 2 _CH 4 + _O 2  _CO 2 + _H 2 O 2 2 3 4 5 2 2 2 2

18. More Examples __ZnS + __O 2  __ZnO + ___SO 2 __C 2 H 6 + __ O 2  __CO 2 + __H 2 O __ NH 3 + __O 2  __ N 2 + __H 2 O