St/Est 404 the Material World

ST/EST 404 THE MATERIAL WORLD CHAPTER 2 NOTES CHAPTER 2 Molecules (pp. 40-50)

1 What is a molecule?

Only a few elements, such as ______and ______, exist in their pure form on Earth. Most atoms combine with atoms of other ______to form ______.

 A molecule is a group of ______or more atoms that are ______bonded together.

Examples: - O2 - ______

- H2O - ______- NaCl - ______

Why do atoms tend to bond with other atoms? ______

 Noble gases (Group VIII) have a ______valence shell; therefore they are extremely ______and rarely ______with other elements.

 Halogens (Group VII) have ______valence electron, so they need to ______one electron to acquire the electron configuration of the nearest noble gas.

1 ST/EST 404 THE MATERIAL WORLD CHAPTER 2 NOTES

 Alkali metals (Group I) have only ______valence electron, so they all tend to ______that electron to resemble a noble gas. Table 2.4: THE TENDENCY OF GROUP A ELEMENTS TO GAIN OR LOSE ELECTRONS Group # IA IIA IIIA IVA VA VIA VIIA VIIIA

Element example Li Be B C N O F Ne

# valence electrons Tendenc y

Octet Rule: The tendency of elements to acquire the configuration of the noble gas ______to them in order to have ______electrons in their valence shell. (Exceptions: Li, ______, and ______acquire the configuration of ______and thus follow the ______rule.)

*Special case: ______- depending on the circumstances it can ______its only electron or it can ______a second electron.

1.1 IONS

In general atoms are electrically ______(equal # of _____ & _____ )

An ion is an atom that has become electrically ______by ______or ______one or more electrons.

ION FORMATION IN METALS

 Since alkali metals (Group I) have only ______valence electron, they all tend to ______that electron when forming ions. When this happens, they acquire a charge of ______.

http://www.nios.ac.in/images/5.1.gif  Since alkaline earth metals (Group II) have ______valence electrons, they all tend to ______electrons when forming ions. When this happens, they acquire a charge of ______.

All metals ______their valence electrons when forming ions and thus form ______charged ions (CATIONS).

ION FORMATION IN NON-METALS

 Since halogens (Group VII) have ______valence electrons, they all tend to ______electron when forming ions. When this happens, they acquire a charge of ______.

Because non-metals all have ______or more valence electrons, they all ______electrons when forming ions and thus form ______charged ions (ANIONS). 1.2 The Nature of Chemical Bonds (EST ONLY) Most atoms, except those of noble gases, have a natural tendency to ______or ______electrons in order to fill their outer shells. When two atoms come together, they will either ______or ______their valence electrons to become ______.

 A ______is the union of two atoms through the ______or ______of one or more electrons.

There are ______main types of chemical bonds: ______bonds and ______bonds.

IONIC BONDS

 An IONIC BOND is usually the result of a transfer of one or more ______from one atom (usually a ______) to another atom (usually a ______). The formation of an ionic bond represented with Lewis structures

F http://www.clickandlearn.org/Gr9_Sci/atoms/bonding.htm

In the Lewis dot diagram above we see that when the sodium atom comes in contact with a chlorine atom, the sodium atom gives up an ______. Both atoms thus acquire an electron configuration similar to that of a ______gas. The sodium atom becomes a

______ion (Na+), and the chlorine atom, a negative ion (Cl-). Since positive and negative charges ______each other, the positive sodium and the negative chloride ion come together to form an______compound.

In the space below, draw a Lewis dot diagram showing the formation of an ionic bond between Magnesium and Bromine (MgBr2).

COVALENT BONDS

Molecular oxygen (O2), ammonia (NH3), and methane (CH4) are examples of the type of bonding where an electron ______reacts with another ______.

 A COVALENT BOND is the result of the ______of one or more electron ______between two ______atoms.

When molecular fluorine (F2) is formed, each atom ______an electron with another fluorine atom so they both have the electron configuration of ______, the nearest Noble gas. In the Lewis structure, the shared electron pair is ______; in the ball and stick model, it is represented by a ______.

 IMPORTANT  F An atom ______share electrons with itself.

F All bonding electrons (unpaired electrons) of one atom must ______with the bonding electrons of the other atom(s). F All atoms (except ____) must have _____ electrons around them in the final diagram. Sometimes two atoms share than one ______pair. In molecular

______(O2), each oxygen atom needs ______more electrons to achieve the configuration of a ______gas, so ______oxygen atoms tend to share two electron pairs. Oxygen atoms are linked in a ______bond. Triple bonds, between atoms of other elements (such as ______) are also possible.

Rutherford-Bohr Atomic Model for O2 Lewis Structure for O2 Ball and Stick Atomic Model for O2

Electron pairs are not always shared ______. Some atoms have a ______force of attraction for electron pairs than others (higher electronegativity). In a ______molecule, the ______atom attracts the electrons more than the two ______atoms do. This causes a certain degree of ______polarity, with the oxygen atom carrying a slightly ______charge, and the hydrogen atoms, a slightly ______one. The unequal covalent bonds are referred to as ______bond.

 Diatomic molecules symbol for a partial charge Some elements do not exist in nature as O individual atoms. Such atoms come in pairs like socks and jeans. They are H diatomic molecules (made up of 2 atoms). To recall which elements are diatomic, H just remember this simple phrase:   “I Have No Bright Or Clever Friends”

1.3 The Rules of Chemical Notation and NomenclatureI2, H (EST2, N2, ONLY)Br2, O2, Cl2, F2

Naming Binary Ionic Compounds

F A binary compound is a compound made up of ______different elements. (bi = ______)

1. Name the metal first 2. Add the suffix –ide to the name of the non-metal

Examples:

NaCl ______

CaF2 ______AgBr ______

Na3N ______KI ______ZnO ______

Mg3P2 ______

Writing Formulas for Binary Ionic Compounds

Use the CROSS-OVER RULE!!!!

The Cross-Over Rule involves writing the charge on each ion as a superscript and then crossing the numbers over and writing them as subscripts without the + and -. Don’t worry, it’s super easy!

F Remember to reduce to lowest terms!

Example: Write the molecular formula for magnesium bromide. Mg2+ Br-

Let’s do a few more:

Aluminum oxide Sodium fluoride Calcium sulfide

______

Stock System for Naming Ionic Compounds Containing Multi-Valent Ions

Some transition metals can form ions with two different charges. Because of the existence of two different ions for these metals, we need a naming system that will enable us to distinguish one from the other. The system we will use is the Stock System. The Stock System involves writing a roman numeral after the name of the metal to distinguish it from its other ion.

Examples:

Fe2+ = ______Fe3+ = ______Pb2+ = ______Pb4+ = ______Cu+ = ______

What do we do when we’re faced with naming this: CuCl2? Is it copper (I) chloride or copper (II) chloride?

We have to do the Cross-Over Rule in reverse! Cu? Cl-

Since the charge on chlorine is ______and there are _____ chloride ions in the formula, that makes 2 × -1 = _____, so the copper has to have a charge of ______in order for the compound to be neutral. So the name of this compound is ______.

Let’s try a few more:

Fe2O3 PbS MnO2 ______

Writing Formulas for Ionic Compounds Containing Multi-Valent Ions

Use the CROSS-OVER RULE just like you do for a regular ionic compound!

Example: Write the molecular formula for chromium (III) chloride. Cr3+ Cl-

Try these:

Nickel (II) bromide ______Gold (III) oxide ______Mercury (II) sulfide ______

Naming Binary Covalent Compounds

F Use prefixes to indicate the number of atoms of each type. F DO NOT use the prefix “mono” on the first element in the formula.

Table 2.18: Prefixes Indicating the Number of Atoms of an Element in a Binary Covalent Compound Number of Atoms Prefix Number of Atoms Prefix One Six Two Seven Three Eight Four Nine Five Ten Examples:

CO2 ______

N2O4 ______CO ______

SF6 ______

PCl3 ______

P4O10 ______

Exceptions to naming covalent compounds

 Some covalent compounds have common names and are not named according to the rule above. (You must memorize the table below!)

Formula Name Formula Name

H2O CH3OH

NH3 C2H5OH

CH4 C6H12O6

C3H8 C12H22O11

C4H10 H2O2

 Hydrogen compounds (ex. HCl, H2S, etc.) DO NOT take prefixes! HCl ______

H2S ______

POLYATOMIC IONS (EST ONLY) p. 44

 A ______ION is a group of ______or more chemically bonded atoms that has become electrically ______by ______or ______one or more electrons.

Table 2.10: Examples of Common Polyatomic Ions Chemical Formula Name Chemical Formula Name - - CH3COO OH Ammonium Nitrate Bicarbonate Nitrite 2- 3- CO3 PO4 2- Chlorate SO4 2- CrO4 Sulfite NOTE: Your teacher will ask you to memorize some of the most common polyatomic ions. Naming ionic compounds containing polyatomic ions (non-binary ionic compounds) Examples:

CaCO3 ______

Mg3PO4 ______NaOH ______

Na2SO4 ______HCN ______

NaHCO3 ______

Writing formulas for ionic compounds containing polyatomic ions (non-binary ionic compounds)

F Use the CROSS-OVER RULE! F If there is more than one polyatomic ion in the formula, you must put brackets around it. F Never change the subscripts of a polyatomic ion! Ex. Ca3(PO4)2 ≠ Ca3P2O8

Example: Write the chemical formula for magnesium hydroxide.

Ammonium nitrate Potassium chromate Aluminum sulfate

______2 Properties of solutions Sometimes atoms and molecules can combine without undergoing a ______reaction to form a ______. Since not ______bonds need to be ______, the different substances that make up a mixture can be ______using physical ______.

 A solution is a ______mixture (consisting of at least one ______and one ______) whose component substances (solids, ______or gases) cannot be ______, even with the aid of a magnifying instrument.  Solute: The component of the substance that is ______in the other. Examples of solutes include salt, sugar, colouring and alcohol.  Solvent: The substance in which the solute ______. Examples of solvents include water, alcohol and acetone.  Aqueous Solution: A solution in which the solvent is ______.

Water is the universal solvent because:  It dissolves many substances. Molecules with ______bonds and molecules with a certain polarity dissolve easily in water. ______molecules, such as oil, rarely dissolve well in water.  It is ______(pH = 7)  ______ Odourless  ______ Doesn’t react

2.1 SOLUBILITY

 SOLUBILITY is the maximum amount of ______that can be ______in a certain volume of ______.

Factors that affect solubility:  ______of the solute  ______of the solvent  ______affects solubility of gaseous solutes  ______(solids tend to become more ______as solvent temperature rises while gases tend to become less ______as solvent temperature rises)

Figure: Solubility of Carbon Dioxide in Water as a Function of Temperature

HTTP://WWW.ENGINEERINGTOOLBOX.COM/GASES-SOLUBILITY-WATER-D_1148.HTML

F SEE Appendix 2 on p. 516 for a list of the solubility (and other characteristic properties) of many common substances.

2.2 CONCENTRATION

 The CONCENTRATION of a solution is the______of ______in a given amount of ______. It is the ratio of the quantity of solute to the quantity of the solution.

DILUTION AND DISSOLUTION The concentration of a solution can be varied in different ways. Change Effect on the concentration

Dilution (______of solvent)

Dissolution (addition of ______)

______(reduction of solvent)

Expressing the Concentration of Aqueous Solutions

1. Concentration: Number of grams of solute per liter of solution (g/L)

Application: What mass of NaOH is needed to prepare 500mL of a 4 g/L NaOH solution?

2. Mass-Volume Percent: Number of grams of solute per100 mL of solution, expressed as a percentage (% m/V)

Application: You have 24g of sugar to prepare a 6%m/V sugar solution. What volume of solution will you make?

3. Volume Percent: Number of millilitres of solute per 100 mL of solution, expressed as a percentage (% V/V)

Application: a) You have 50mL to prepare a 6%V/V alcohol solution. What volume of solution will you make?

b) You add 75 mL of acetone to 1205mL of water. What is the concentration of the solution in %V/V?

4. Mass Percent: Number of grams of solute per 100 mL of solution, expressed as a percentage (%m/m)

Application: What is the mass of NaCl in 400g of an 8%m/m brine solution?

CONCENTRATION IN PPM

When the amount of solute in the solution is very small, the concentration can be expressed in ______.

 The CONCENTRATION in PPM (“ ______”) is the number of parts of solute in a ______parts of solution. 1g 1mg 1 ppm = = = 1 mg/L 1000000g 1000g

VERY IMPORTANT!!!!  X%(m/V) = X g/100mL  Example: 5%(m/V) = 5g/100mL = 5000g/0.1L = 50 000 mg/L or 50 000 ppm 15 ST/EST 404 THE MATERIAL WORLD CHAPTER 2 NOTES

MOLAR CONCENTRATION (AKA Molarity) (EST ONLY)  Expressed as moles per liter (______)

 MOLAR CONCENTRATION corresponds to the number of ______of dissolved solute particles in a ______of solution.  Molar concentration is symbolized by placing the ______formula for the measured substance inside ______brackets. Example [NaCl] = 0.5 mol/L means the ______concentration of sodium chloride solution equals ______mol/L.

Molar concentration of a solution can be calculated using the formula:

The units for molarity are mol/L or M (in this case, M or MC stand for molar concentration, not molar mass)

Example: Suppose 58.5 g of sodium chloride are dissolved in 500 mL of solution. Calculate the molar concentration of this solution following the method illustrated on page 54.

More Practice

a) If 20g of KNO3 is dissolved in enough water to make 500mL of solution, what is the molar concentration of the solution? (Answer 0.4M)

b) What mass of CaF2 is needed to prepare 250mL of a 0.1mol/L solution? (Answer 1.95 g)

F The following is not covered in Chapter 2 of OBSERVATORY. It is, however, important information that you need to know for your lab exam in June.

HOW TO PREPARE A SOLUTION

How would you prepare 250mL of a 20 g/L solution of cobalt (II) chloride, CoCl2?

1. Calculate the mass of solute needed.

2. Using an electronic balance weigh out the mass of the solute needed. 3. Pour a small amount of water into the volumetric flask. 4. Pour the solute into the flask. 5. Add water up to etched line. Use a pipette near the end (bottom of meniscus should be on the etched line) 6. Cap the flask, invert the flask and shake. Repeat 3 times.

DILUTING SOLUTIONS

To dilute a solution is simply to add ______to a more concentrated solution. Even after you do this, the amount of ______is the same in both the original concentrated solution and the new diluted solution.

mc = mass of solute in concentrated solution Recall that C = m/V, so m = C  V md =mass of solute in diluted solution mc = md

DILUTION FORMULA Cc Vc = Cd Vd

Cc = concentration of the concentrated solution Cd = concentration of the diluted solution (in %m/m, %V/V, %m/V, or g/L) (in %m/m, %V/V, %m/V, or g/L)

Vc = volume of concentrated solution Vd = volume of diluted solution (in mL or L) (in mL or L)

Some tips to remember:

Cc  Cd (ex. 30%V/V  6%V/V) Vd = Vc + Vwater added

Vc  Vd (ex. 5 mL  25 mL) Vwater added = Vd - Vc

Example 1: You have a bottle of 5%m/V bleach solution. You need to prepare 500 mL of a 2% m/V bleach solution. What volume of concentrated bleach will you use?

Example 2: What is the final concentration when 100 mL of water is added to 200 mL of a 12 g/L salt solution?

Example 3: What volume of water must be added to 500 mL of a 30%V/V hydrogen peroxide solution to dilute it to 6%V/V?

Example 4: Using 300 mL of a 45 g/L solution of lithium chloride, LiCl, a student must prepare a 15 g/L solution of lithium chloride. What is the volume of the resulting solution?

2.3 ELECTRICAL CONDUCTIVITY

Pure water does not ______electricity. How then does it does carry and electric current? It is the substances ______in the water that conducts the electricity.

 An ______is a substance that when dissolved in water, allows an ______to flow through the solution.

 The ELECTRICAL ______of a solution is a measure of its ability to allow ______to flow through it.

 A ______is a substance that is soluble in water but DOES NOT conduct electricity.

ELECTROLYTIC DISSOCIATION

 When an ______is dissolved in water, it separates into two ______of opposite charge, one ______and one ______.  This separation, known as ______, is a ______change.

The chemical equation for the electrolytic dissociation of sodium chloride is:

F The H2O over the arrow indicates that the change takes place when the ______is placed in water. F The ions formed during this process ______electricity.

A non-electrolyte does not conduct electricity when dissolved in water because it does not ______.

F All covalent compounds (except acids) are non-electrolytes.

Here’s what happens when you dissolve methanol (CH3OH) in water:

THE STRENGTH OF ELECTROLYTES (EST ONLY)

To determine whether a solute is an electrolyte:

1. Dissolve it in ______2. Place two ______in the solution 3. Connect to a power supply and a light bulb 4. Does the bulb light up?

Strong Electrolytes Weak Electrolytes Non-Electrolyte Substances that dissociate Substances that only Substances that ______(100%) ______dissociate produce ions when dissolved in when dissolved in water. when dissolved in water. water. + - 2+ 2- NaCl –H2O→ Na (aq) + Cl (aq) CaCO (s) –>50% → Ca (aq) + CO3 (aq) C12H22O11(s) –H2O→ C12H22O11(aq) 100 molecules 100 ions + 100 ions 100 molecules 50 ions + 50 ions 100 molecules 100 molecules

no light

 The strength of an ______is the degree to which it ______into ions. The higher the degree of dissociation, the ______the electrolyte.

TYPES OF ELECTROLYTES Acids Bases  Found in fruit juices, ______ Found in many ______

______, & gastric juices products and in some  pH ______than 7 ______ Taste ______medication. Blood and ______water  Turns blue litmus paper ______are also slightly basic.  Neutralize ______ pH ______than 7  Release H+ ______when dissolved in  Taste ______water  Turn ______litmus paper blue  Molecular formula often begins with  Neutralize ______the symbol for a ______ Feel ______to the touch atom followed by a nonmetal (HCl,  Dissolve ______and oils

- HNO3, HSO4, HF, H2CO3)  Release ______ions (OH

F Acetic acid (aq)) when dissolved in water (______) is  Molecular formulas begins with a metal the exception to this rule. and usually ends in “OH”:

Examples of acid solutions: NaOH, KOH, NH4OH, Ca(OH)2, + - 1. Hydrochloric acid: HCl → H (aq) + Cl (aq) Al(OH)3 + - 2. Nitric acid: HNO3 → H (aq) + NO3 (aq) + 2- F Exceptions: Alcohols! CH3OH, 3. Sulfuric acid: H2SO4 → 2H (aq) + SO4 (aq)

C2H5OH  these are NOT bases and

NH3 is a base yet it does not end in OH Examples of basic solutions:

+ - 1. Sodium hydroxide: NaOH → Na (aq) + OH (aq) + - 2. Potassium hydroxide: KOH → K (aq) + OH (aq) 2+ 3. Magnesium hydroxide: Mg(OH)2 → Mg (aq) + - 2OH (aq)

 Constitute a ______class of substances that figures ______in human diets

 Most are made up of a ______and one or more ______.  pH = ______(______)  Have ______effect on litmus paper

 A SALT is a substance produced by the ______bonding of a ______ion and a ______ion (other than H+ and OH- ions). Examples of salt solutions:

1. sodium chloride: NaCl,

2. potassium bromide: KBr

3. calcium chloride: CaCl2

4. silver nitrate: AgNO3

F Not all salts dissolve easily in water.

2.4 pH Acidic, basic and neutral solutions can be distinguished by their ______. A solutions pH can be measured using a pH ______or a pH ______.

THE pH SCALE  ranges from ______to ______ if the pH the solution is ______ if the pH = 7, the solution is ______ if the pH the solution is ______ the pH scale is ______, which means that a difference of one unit between two substances actually indicates that one of the substances is ______times more acidic than the other.

Examples: a) What is the pH of an acetic acid solution that is 100 times more diluted than an acetic acid solution with a pH of 2?

______

b) What is the pH of an ammonia solution that is 1000 times more concentrated than an ammonia solution with a pH of 10?

______

Figure 1: The pH of some common substances http://islandwood.org/kids/stream_health/Data/pH_scale.jpg

State whether the following substances are acidic, basic, or neutral:

Susbtance Acidic, Basic, Neutral? Vinegar Bleach Coke Detergent Sea Water MORE ON pH (EST ONLY)

The pH of a solution is actually an indication of the concentration of ______ions END(_____) OF in that CHAPTER solution. 2 NOTES FOR ST

How is pH related to hydrogen ion concentration?

+ pH = -log [H ] where [H+] = hydrogen ion concentration in mol/L

So, if the hydrogen ion concentration in a solution is 1 10-3 mol/L, then the pH = ______.

+ -pH Likewise, [H ] = 10

So, if the pH of a solution is 8, then the hydrogen ion concentration is ______.

F See Table 2.30 on page 61 of your textbook.

THE pOH SCALE

The complete opposite of the pH scale, it communicates hydroxide ion concentrations, [OH-], in a wide variety of substances.

- pOH = -log [OH ] where [OH-] = hydroxide ion concentration in mol/L

So if the hydroxide ion concentration in a solution is 1 10-6 mol/L, then the pOH = ______.

- -pOH Likewise, [OH ] = 10

So, if the pOH of a solution is 10, then the hydroxide ion concentration is ______.

How are pH and pOH related? pH + pOH = 14

So if the pH of a solution is 5, then the pOH is ______.

END OF CHAPTER 2 NOTES FOR EST