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Chapter 4

Chapter 4. Chemical Reactions. 4.1 Intro to Chemical Reactions. Chemical Reactions. Are the cause of chemical changes Remember a chemical change causes one substance to change into another!. Physical vs. Chemical Change Review. Physical change changes things like shape & size

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Chapter 4

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  1. Chapter 4 Chemical Reactions

  2. 4.1 Intro to Chemical Reactions

  3. Chemical Reactions Are the cause of chemical changes Remember a chemical change causes one substance to change into another!

  4. Physical vs. Chemical Change Review Physical change changes things like shape & size Also physical changes… • Stirring • Dissolving • Drying • Separating Remember physical changes are REVERSIBLE

  5. Physical vs. Chemical Change Review Chemical Change changes the actual structure of the molecule turning it into a completely different substance Chemical Changes are IRREVERSIBLE- they cannot be undone If they can be undone the substance has to go through another chemical change to do so

  6. Test your memory… What are the 4 signs of a chemical change?

  7. Chemical Bonds and Intermolecular Forces The molecules in the substance are held together by much weaker forces called Intermolecular Forces The atoms in a molecule are held together by super strong forces called Chemical Bonds

  8. Energy Requirements Both physical and chemical changes require energy A Physical change only breaks intermolecular forces, while a chemical change has to break the chemical bonds Which kind of change do you think requires more energy?

  9. Chemical Bonds We just learned that molecules are created and held together by chemical bonds Chemical bonds form when the electrons from one atom are either SHARED or TRANSFERRED with another atom Reminder: Electrons are the negatively charged particles that orbit the nucleus

  10. Chemical Bonds There are 2 types of chemical bonds: Ionic and Covalent Ionic Bonds: Electrons are transferred from a metal to a non-metal Covalent Bonds: Electrons are shared between two non-metals

  11. Draw a “stairs” on your periodic tableThe elements to the left = metalsThe elements to the right = non-metals

  12. Practice Decide whether the element is a metal or non-metal Na Au P O Ne Cu Mg H Pb

  13. Practice Is the molecule held together by an ionic bond or a covalent bond? CH4 MgO KBr CuCl2 HI NH3

  14. Reactivity Have you noticed I use a lot of the same elements again and again? This is because certain elements are more reactive than others!

  15. 4.2 Understanding and Balancing Chemical Equations

  16. Chemical Equations Chemical equations are written to show the chemical reaction 3 major parts to a chemical equation: • Reactant • Product • Yield Arrow

  17. Chemical Equations Reactants NaCl + AgNO3→ NaNO3 + AgCl End results, always on the right Products Starting materials, always on the left Yield Arrow shows the direction of the reaction, means “reacts to form”

  18. Balancing Equations When you have a chemical equation you need to make sure it is balanced Law of Conservation of Mass: Mass cannot be created or destroyed So the total mass and the number of atoms of each element must be the same on both sides of the equation!

  19. Balancing Equations coefficient We balance equations by adding a coefficient in front of a molecule to change the number of atoms in the equation We NEVER change the subscripts, this changes the molecule into something totally different! subscript

  20. Coefficients Coefficients get distributed throughout the entire molecule • 4 H’s and 2 O’s The coefficient also tells you the number of moles of the molecule it is in front of • 2 moles of H2O

  21. How to balance equations There is no “right” way to balance an equation but most people find this general process the easiest way to do so Most of balancing equations is just trial and error • Start out listing the number of atoms of each element in the equation • Look for any unbalanced atoms and add a coefficient in front of the molecule that would balance that atom • Repeat step 2 until you have a balanced equation

  22. Balancing Equations Example: H2 + O2 H2O Step One:ReactantsProducts 2 H 2 2 O 1 (x2) Step Two:H2 + O2 2H2O ReactantsProducts (x2) 2 H4 2 O2 Step Three: 2H2 + O2 2H2O

  23. Try it on your own!Zn + HCl → ZnCl2+ H2 Step One:ReactantsProducts Zn H Cl Step Two:Zn + HCl → ZnCl2+ H2ReactantsProducts Zn H Cl

  24. One more… CH4+ O2→ CO2+ H2O

  25. Conservation of Energy We already know that matter cannot be created or destroyed, the same is true for energy Conservation of Energy: Energy cannot be created or destroyed, it can only be transferred or transformed

  26. Energy and Chemical Reactions In a reaction the chemical bonds are broken and then new bonds are formed. Because of this reactions either give off or gain energy! Energy usually = heat but can also mean light or sound

  27. Endothermic &Exothermic Reactions Endothermic Reactions: Need to absorb energy to go from reactants to products • Products have higher energy than the reactants • Takes more energy to form new bonds then to break the original bonds Exothermic Reactions: Releases energy to go from reactants to products • Products have lower energy than the reactants • Takes more energy to break the bonds then to form new ones

  28. Conservation of Energy What happens to energy in a chemical reaction The reaction = the “system” The room = the “surroundings” 1. The energy of the system + surroundings stays the same. 2. Energy gained by the system must be lost by the surroundings. 3. Energy lost by the system must be gained by the surroundings.

  29. Exothermic Reactions Exothermic reactions give off energy, so why don’t they happen spontaneously? Even though they give off a lot of energy, they require a small amount of energy to be added to the system so the reaction can start After this small amount of energy is added then a HUGE amount of energy is released This small amount of starter energy is called activation energy

  30. 4.3 Types of Chemical Reactions

  31. Solutions Remember that in a solution the solute is dissolved in the solvent When the solute dissolves the molecule completely breaks down and freely floats around Reactions can only occur if molecules can move around and touch each other so MANY reactions require the molecules to be dissolved in a solution

  32. Aqueous Solutions Any solute dissolved in water is called an aqueous solution This is so important and so common that we consider being dissolved in water to be like a fourth phase (solid, liquid, gas, aqueous)

  33. Writing phase of matter in chemical reactions When we write a chemical equation we need to include the phase of each reactant/product in the equation We use the symbols (s), (l), (g), and (aq) to show what phase the reactants and products are in: (s) = solid. (l) = liquid. (g) = gas. (aq) = dissolved in water.

  34. Precipitate Reactions BaCl2(aq) + Na2SO4(aq) → BaSO4(s) + 2NaCl(aq)  When two aqueous substances react and form a solid substance a precipitatereaction has formed A precipitate forms when a molecule cannot dissolve in water A precipitate reaction is a type of a double replacement reaction

  35. Synthesis and Decomposition Synthesis- Two or more molecules coming together to form one new molecule H2 + O2 H2O Decomposition- One molecule breaks down to form two or more new molecules H2CO3  →   H2O + CO

  36. Replacement Reactions Single Replacement: One element trades places with another element in a molecule. Example: Zn + HCl → ZnCl2+ H2 Double Replacement: This is when the two parts of two molecules switch places, forming two completely new molecules Example: NaCl + AgNO3 → NaNO3 + AgCl

  37. Acid/Base Reaction

  38. Acid/Base Reactions An acid/base reaction will ALWAYS produce water an a salt Salt is NOT only the salt we use on food, it is any substance formed from the positive ion from a base and the negative ion from an acid HBr+ NaOH → NaBr + H2O

  39. Combustion Reaction Reactants: Hydrocarbon (only C’s and H’s) and Oxygen react and always form Products: Water(H2O) and Carbon Dioxide (CO2) CH4 + O2 → CO2 + H2O

  40. Practice Match the chemical reaction to its type of reaction • Acid/BaseNaCl(aq) + AgNO3(aq) → NaNO3(aq) + AgCl(s) • Combustion Mg + 2H2O →Mg(OH)2 + H2 • Decomposition8 Fe + S8  →8 FeS • Double Replacement 2C2H6+ 7O2→ 4CO2 + 6H2O • Precipitation 2 H2O →2 H2 + O2 • Single Replacement MgO + 2KCl → K2O + MgCl2 • Synthesis HBr+ NaOH →NaBr + H2O

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