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General Chemistry 1 > Chemical Reactions Definitions and Formulas > Flashcards

Chemical Reactions Definitions and Formulas Flashcards

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1
Q

Ions in Aqueous Solutions

A
  • Pure water does not conduct electricity
  • Water can dissolve many things, referred to as the “Universal Solvent”
  • Reactions taking place in water are called aqueous solutions.
  • water becomes a conductor of electricity due to the production of ions
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2
Q

Who is Svante Arrhenius?

A
  • Swedish Chemist
  • Proposed the idea that some substances dissociate into cations and anions when dissolved in water.
  • His ideas are now known as the theory of electrolytic dissociation.
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3
Q

Electrolyte

A

Substance that dissolves in water to give an electrically conducting solution.

e.g NaCl(s) → Na+(aq) + Cl-(aq)

Most ionic compounds are strong electrolytes, but there are a few exceptions.

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4
Q

Non-Electrolytes

A

Substance that dissolves in water to give a non-conducting or very poorly conducting solution.

e.g. C12H22O11(s) + H2O(l) → C12H22O11(aq)

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5
Q

How do you determine the strength of an electrolyte?

A

The strength of an electrolye depends on the extent of the dissociation or ionization in the solution.

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6
Q

Dissociation

A

Seperation or splits apart as the salt dissolves

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7
Q

Ionization

A

Formation of ions

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8
Q

Strong electrolyte

A

an electrolyte that exists in solution almost entirely as ions

dissociates to a large extent; 70 to 100%

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9
Q

Weak Electrolyte

A

dissolves in H2O to give a small percentage of ions, about 1%

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10
Q

Summary of Electrolytes

A
  • Most ionic compounds are strong electrolye with a few exceptions
  • Few molecular compounds are strong electrolytes
  • Most molecular compounds are weak electrolytes or non-electrolytes
  • Most organic compounds are molecular and non-electrolytes; however, carboxylic acids and amines are week electrolytes.
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11
Q

Solubility

A
  • Ability to dissolve in water varies according to the substance.
  • Some substances are very soluble while others are insoluble.
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12
Q

Describe reactions involving ions.

A
  • Molecular equation
  • Complete or total ionic equation
  • Net Ionic equation
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13
Q

Molecular equation

A

Chemical equation in which the reactants and products are written as if they were molecular substances, even though they may acutually exist in solution as ions.

Pb(NO₃)₂(aq) + 2K I(aq) 2KNO₃(aq) + PbI₂(s)

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14
Q

Complete or total ionic equation

A

Chemical equation in which strong electrolytes such as soluble ionic compounds are written as seperate ions in the solution.

Pb²⁺(aq) + 2NO₃⁻(aq) + 2K ⁺(aq) + 2I⁻(aq) → 2K⁺(aq) + 2NO₃⁻(aq) + PbI₂(s)

Soluble solutions have (aq) subscripts and insoluble solid substances have (s) subscripts.

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15
Q

Spectator Ions

A
  • Ions the do not undergo any change in the reaction and appear on both sides of the equation.
  • They are there to balance the charge
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16
Q

Net ionic equation

A
  • An ionic equation from which spectator ions have been canceled.
  • Only the ions undergoing change are shown

—Pb²⁺(aq) + 2NO₃⁻(aq) + 2K ⁺(aq) + 2I⁻(aq) → 2K⁺(aq) + 2NO₃⁻(aq) + PbI₂(s)

to

Pb²⁺(aq) + 2I⁻(aq) → PbI₂(s)

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17
Q

Rules for writing Net Ionic Equations

A
  1. Write strong electrolytes in their ionic form
  2. Write weak electrolytes in molecular form
  3. Write non-electrolytes in molecular form
  4. Insoluble substances, solids or precipitates, and gases in their molecular form
  5. Should only have substances that have undergone a chemical change.
  6. Equation must be balances
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18
Q

What are the 3 types of chemical reactions

A

There are three types:

  1. Precipitaton reaction
  2. Acid-base reeaction
  3. Oxidation - Reduction reaction
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19
Q

Precipitation Reaction

A
  • When you mix two ionic substances and a solid precipitate result.
  • Soluble reactants yield an insoluble product that drops out of the solution.
  • The driving force for the reaction is the formation of the stable solid product which removes material from the aqueoaus solution
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20
Q

Acid-Base reactions

A

An acid substance that involve the transfer of a proton (H+)

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21
Q

Oxidation-Reduction reaction

A

Reaction that involves the transfer of an electron.

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22
Q

Precipitate

A
  • An insoluble solid compound formed durinf a chemical reaction in solution.
  • To predict whether a precipitate will form on mixing aqueous solutions of two substances, you need to know the solubility of each potential product
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23
Q

Solubility

A
  1. How much of each compound will dissolve in given amount of solvent at a given temperature.
  2. Low solubility in water
    • it is likely to precipitate from an aqueous solution
  3. Hight solubility in water
    • no precipitate will form
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24
Q

Exchange reaction

A
  • Also known as metahesis reaction.
  • reaction between two compounds that, when written as a molecular equation, appears to involve the exchange of parts between the two reactants.
  • The two cations exchange partners
    • AB + CD → CB + AD
  • Note: if an insoluble precipitate forms, then reaction is possible, if no precipitate, then no reaction.
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25
Q

Acid-Base Reaction

A
  • Important class of chemicals
  • acids:
    • have a sour taste
    • No feel
    • litmus change from blue to red
  • bases:
    • bitter taste
    • soapy feel.
    • litmus change from red to blue
26
Q

Acid-Base indicator

A
  • A dye used to distinguish between acidic and basic solutions by means of the color changes it undergoes in these solutions.
  • Note: phenolphthalein sn indicator which is colorless in acid and pink in basic solution.
27
Q

pH value example

A

28
Q

Arrhenius Acid Definition

A

A substance that produces H+ when dissolved in water

29
Q

Arrenius base definition

A

Substance that produces hydroxide ions, OH- when dissolved in water.

30
Q

Problems with Arrhenius acid-base definition

A
  1. Reactions has to be in water
  2. Some bases, such as NH3, ammonia and amines do not have hydroxides in their formula.
31
Q

Bronsted-Lowry Acid Definition

A
  • the molecule or ion that donates a proton to another molecule or ion in a proton-transfer reaction
  • A proton H+ donor
32
Q

Bronsted-Lowry Base Definition

A
  • Molecule or ion that accepts a proton in a proton-transfer reaction.
  • a H+ acceptor
33
Q

Bronsted-Lowry acid-base concept

A

​

  • Note that the acid becomes a base because it is lacking a proton and the base becomes an acid because it aquired a proton.
  • For simplicity, H+ is usually written but it is really H3O+, called a hydronium ion.
34
Q

Hydronium Ion

A
  • it is nothing more than a bare proton and does not exist by itself in aqueous solution
  • in water, it combines with a polar water molecule to form a hydrated hydrogen ion
  • the H+ is attached to to a polar water molecule forming a bond with one of the two pairs of unshared electrons.
35
Q

Neutralization Reaction

A
  • Reaction of an acid and base to form an ionic compound (salt) and possibly water
  • Driving force of a neutrilization reaction is the ability of H+ ion and an OH- ion to react to form a molecule of unionized water.
  • If the spectator ions, in a neutralization reaction are soluble salts,they remain in the solution.
  • Evaporation of the solution to dryness yields the pure salts.
36
Q

What is the exception to the neutralization Reaction?

A
  • Water is produced in most neutralization reactions
  • except for when the reaction of an acid is with a weak base (ammonia)
37
Q

Monoprotic acids

A

one acidic hydrogen atom per acid

e.g. HCl, HNO3, HCN, HBr,HI, HClO4

38
Q

Polyprotic acids

A
  • two or more acidic hydrogens per molecule.
  • e.g. H₂CO₃, H₃PO₄, H₂SO₄,etc .
  • They can form a series of salts with different amounts of base.
  • Some of the salts formed have hydrogen atoms so they are acidic salts and can undergo neutralization with bases.

H₃PO₄(aq) + NaOH → NaH₂PO₄( aq) + H₂O(l)
H₃PO₄(aq) + 2NaOH → Na₂HPO₄( aq) + 2H₂O(l)
H₃PO₄(aq) + 3NaOH → Na₃PO₄( aq) + 3H₂O(l)

  • So salts, NaH₂PO₄ and Na₂HPO₄ are acidic salts.
39
Q

Acid-Base reactions with Gas Formation

A
  • Some salts like carbonates, sulfites, and sulfides react with acid to produce gas.
  • Baking soda (NaHCO₃) reacts with an acid to produce carbonic acid, H₂CO₃, which decomposes to CO₂ and water.
  • The net ionic equations for some gas forming reactions are :
    • HCO₃⁻ + H⁺ →CO₂(g) + H₂O(l)
    • CO₃²⁻ + 2H⁺ → CO₂(g) + H₂O(l)
    • HSO₃⁻ + H⁺ → SO₂(g) + H₂O(l)
    • SO₃²⁻ + 2H⁺ → SO₂(g) + H₂O(l)
    • S²⁻ + 2H⁺ → H₂S(g)
    • HS⁻ + H⁺ → H₂S(g)
40
Q

Oxidation-reduction reactions

A
  • Transfer of electrons, known as redox reactions
  • Oxidation
    • Loss of electrons (or increase of oxidation number)
  • Reduction
    • Gain of electrons (or decrease of oxidation number)
  • Mnenomic
    • LEO the lion goes GER
      • LEO: Lose electrons - oxidation
      • GER: Gain electrons - reduction
    • OIL RIG
      • OIL: Oxidation is loss of electrons
      • RIG: Reduction is gain of electrons
41
Q

Oxidation Number

A
  • Actual charge on a monatomic ion or a hypothetical charge assigned to the atom in the substance by simple rules.
  • It is a way to keep track of electrons in a redox reaction
  • by comparing the oxidation number of an atom before and after reaction, we can tell whether it has gain or lost electrons.
  • For the oxidation number, the plus or minus sign is in front of the number to differentiate it from the electronic charge.
42
Q

Rules for assigning oxidation numbers

A
  • Elements
    • is zero
  • Monatomic ions
    • equal to the charge on it
  • Oxygen
    • most compounds is -2, except for H2O2 where it is -1
  • Hydrogen
    • most compounds is 1, except in binary compounds with a metal then it is -1
  • Halogens
    • Flourine is -1
    • Cl,Br,I is -1 in binary compounds, except when combined with a halogen above it or the other element is oxygen
  • Compounds and Ions
    • Algebraic sum in compounds is zero
    • Algebraic sum in a polyatomic ion is equal to the charge on the ion.
43
Q

Describing Oxidation-Reduction Reactions

A
  • They occur simultaneously.
  • A half reaction tells you which is the oxidation and which is the reduction reaction.
  • It is possible to identify the reducing and oxidizing agent by the half reaction.
44
Q

Reducing Agent

A
  • Causes reduction
  • loses one or more electrons
  • undergoes oxidation
  • oxidation number of atom increases
45
Q

Oxidizing Agent

A
  • Causes oxidation
  • Gains one or more electrons
  • undergoes reduction
  • oxidation number of atom decreases
46
Q

Oxidizing and Reducing Agent

A
  • The reducing agent is itself oxidized when it gives up electrons
  • The oxidizing agent is itself reduced when it accepts electrons.

47
Q

What are the four common oxidation-reduction reactions?

A
  1. Combination reaction
  2. Decomposition reaction
  3. Displacement reaction
  4. Combustion reaction
48
Q

Combination Reaction

A

One of the oxidation-reduction reactions

Two substances combine to form a third:

2H2 + O2 → 2H2O

49
Q

Decomposition reaction

A

One of the oxidation-reduction reactions:

One compound decomposes to give several substances

2HgO(s) → 2Hg(l) + O2(g)

50
Q

Displacement reaction

A
  • One of the oxidation-reduction reactions
  • An element displaces another element from a compound
  • Note: a metal will displace from solution the ions of any metal that lies below it in the activity series
  • e.g. a thermite reaction

51
Q

Combustion Reaction

A
  • One of the oxidation-reduction reactions
  • reaction with oxygen, with the rapid release of heat produces a flame.
  • CH4(g) + 2O2(g) → CO2(g) + 2H20(g)
  • Oxygen changes oxidation number from 0 to -2
52
Q

Balancing Simple Oxidation-Reduction reaction

A
  • Charges must be balanced in as well as each element.
  • total increase in oxidation number equals total decrease in oxidation number
53
Q

Half-Reaction Method

A
  • Identify oxidation/reduction half reactions by assigning oxidation numbers
  • Balance charge by adding electrons to the more positive side.
    • Right (product) side for the oxidation half reaction
    • Left (reactant) side for the reduction half reaction
  • Multiply by a factor to make the number of electrons equal. Cancel the electrons.

Example:

Mg(s) + N2(g) → Mg3N2(s)

0 0 +2 -3

Balance oxidation half-reaction

Mg(s) → Mg2+ + 2e-

N2(g) + 6e- → 2N3-

Need to multiply each half-reaction by a factor that will cancel the electrons

3(Mg(s) → Mg2+ + 2e-)

1(N2(g) + 6e- → 2N3-)

3Mg(s) + N2(g) + 6e- → 3Mg2+ + 2N3- + 6e-

Get rid of the electrons:

3Mg(s) + N2(g) → 3Mg2+ + 2N3-

3Mg(s) + N2(g) → Mg3N2(s)

54
Q

Solution

A
  • Reactions in a solution in faster
  • Parts of a Solution
    • Solute: substance dissolved or less abundant
    • Solvent: dissolving agent or more abundant
    • Concentrion: quantity of solute in a standard quantity of solution
      • Dilute: solute concentration is low
      • Concentrate: solute concentration is high
55
Q

Molar concentration

A
  • also called Molarity (M)
  • is moles of solute dissolved in one liter of solution
  • Volume measurements are more convenient than mass measurements. Molarity can be used as a conversion factor.
56
Q

Dilution of Solutions

A
  • Process of preparing a more dilute solution by adding solvent to a more concentrated one.
  • The addition of solvent does not change the amount of solute in a solution, but does change the solution concentration.
  • Moles of solute = molarity x Liters of solution
  • Formula:
    • Minitial x Vinitial = Mfinal x Mfinal
57
Q

Quantitative analysis

A

Determination of the amount of a substance or species present in a material.

58
Q

Gravimetric Analysis

A

a type of analysis in which the amount of a species in a material is determined by converting the species to a product that can be isolated completely and weighed. Precipitation reactions are used in gravimetric analysis.

59
Q

Volumetric analysis

A
  • Determining the amount of a substance by using the volume of another substance of know concentration required for a complete reaction.
  • Measures the volume of one reagent required to react with a measured mass or volume of another reagent.
  • based on titration.
60
Q

Titration

A

determining the concentration of a solution by allowing a carefully measured volume to react with a standard solution of another substance, whose concentration is known.

General Chemistry 1 (19 decks)

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