ICPS Spring Final (Unit 5-12)

Objectives: 

• Coulomb's Law — describes the force between two charged particles

  • Calculate using Coulomb's Law: force between two charged particles in simple scenarios

  • Explain using Coulomb’s Law: real-world observations of simple electrostatic scenarios 

• Explain: force between charged particles increases as the charges increase (Use the equation)

• Explain: force between charged particles decreases as their distance increases (Use the equation)

• Attractive (pulling together) v.s. repulsive (pushing apart) forces based on particles’ charges

Coulomb’s Law

Electrostatic Force (Fe) = kQ1 Q2d2

Q = charges of the objects | d = distance between objects | k = Coulomb's constant

• Equation calculates & describes the electrostatic force between charged objects

  • Fe < 0 → charges attract

  • Fe > 0 → charges repel 

  • | Fe | → magnitude of electrostatic force 

• Units

  • Charge: Coulombs (C)

  • Distance: Meters (m)

  • Constant: Nm2C2

  • Force: Newtons (N)

• Similar to the equation for gravitational force: GM1 M2d2

Conclusion: 

The stronger the charge on each object, the stronger the interaction. 

Force is directly related to the product of charges. 

The closer the objects are to each other, the greater the attraction/repulsion. 

Force is inversely related to the distance. 

• Calculate electrostatic force using Coulomb’s Law by plugging in the charges, distance, and constant into the equation

Electric Charge

Basic property of matter responsible for attraction & repulsion between charged particles.

• Objects with opposite electrical charges attract while same charges repel 

• Electrons (-) & protons (+) have opposite charges that are same in magnitude

• Most matter is neutral (same no. protons & electrons)

• Total charge of a system is always preserved, but it can be transferred from object to object (similar to conservation of energy)

Real-World Scenario Examples

1. Two strips of tape. Stick one piece on a table and the other on top of the first piece. Rub the strips together, then pull off the strips one at a time. Hold the strips apart and slowly bring them together. The two strips attract and move towards each other. 

Explanation: Rubbing one strip of tape onto another transfers electrons from the top strip to the bottom, making the bottom tape negative and the top tape positive. Opposite charges attract, thus, the tapes are attracted to each other. 

2. Rub a balloon on your hair, then bring it close to a wall, placing it on the vertical surface. The balloon will stick to the wall. 

Explanation: Hair is a medium that holds a lot of electrons. Rubbing the balloon against hair transfers electrons from the hair to the balloon, making the balloon negatively charged. When the balloon is brought close to the neutral wall, its electrons are repelled by the balloon’s electrons, leaving a region of positive charge which the negative balloon is attracted to. 

Objectives: 

Matter

• Define: elements, compounds, molecules, atoms, mixtures

• Homogeneous v.s. heterogeneous mixtures

• Understand: components of a mixture retain their individual properties

• Classify different types of matter based on their composition

Atomic Structure

• Define: atoms are the basic building blocks of matter

• Identify & Describe: three main subatomic particles—protons, neutrons, electrons

  • Relative masses, charge, location

• Explain: the atomic nucleus concept & its composition

• Understand: the concept of an electron cloud or energy levels 

• Models of the atom & the scientific findings that created them

• Understand: atomic number & element symbols (relation to number of protons)

Isotopes & Ions

• Define: Isotopes & Ions

• Understand: mass number (relation to protons & neutrons) 

• Cations (positive ions) v.s. Anions (negative ions)

Atomic Mass

• Define: atomic mass

• Understand: how to calculate average atomic mass of an element 

  • Given relative abundance of its isotopes

Matter

Atoms: smallest unit of matter that still has all of the properties of its element

Molecules: structure or particle that contains multiple atoms bonded together

Pure Substances: 

Elements — Substance composed of only 1 element from the periodic table

• Ex. helium, oxygen

Compounds — Substance composed of one type of particle, but the particle is composed of multiple elements 

• Ex. pure water, table salt

Mixtures:

• Composed of multiple types of particles (typically can be separated)

• Components of a mixture retain their individual properties (creating mixtures does not involve chemical reactions or bond formations)

Homogeneous Mixture — Uniform throughout (mixture has the same composition everywhere)

• Ex. cup of coffee, mixed lemonade, brass

Heterogeneous Mixture — Not uniform throughout (the ratio of particles is different in different places)

• Ex. layered cake, chocolate chip cookies, orange juice with pulp

Atomic Structure

Atoms are the building blocks of matter. 

Subatomic Particles:

Additional Information: 

Protons & Neutrons

• About the same size (1 atomic mass unit)

• Nuclear charge (composed of proton & neutron charges) acts altogether as one charge 

Electrons

• Smallest particle 

• Always moving 

• Behaves as an individual particle with its own charge (each has -1 charge)

Atomic Nucleus 

• Small, dense region at the center of an atom consisting of protons & neutrons

  • Positively charged 

• Extremely small volume, but makes up nearly the entire mass of the atom

• Held together by the strong nuclear force (strongest fundamental force) 

• Discovered in 1911 through Rutherford’s Gold Foil Experiment 

Electron Cloud

• Area around the nucleus of an atom where electrons are most likely to be found

• Quantum model: says that we cannot know exactly where an electron is at any given time, but they are likely to be in the electron clouds

• Orbitals are the areas with a dense electron cloud 

  • Dense cloud = high likelihood of finding an electron there

Models of the Atom & Their Scientific Findings

Element Symbols 

Atomic Number:     Number of protons in the atom 

• Determines the element 

• Neutral atoms: no. of electrons = no. of protons

Isotopes

Atoms of an element with a different number of neutrons. 

Isotope number = Mass Number of the atom

• Sum of the # of protons + # of neutrons in the atom

• Differs from Atomic Mass

• Isotopes depend on the # of neutrons 

  • # protons cannot change, as that would change the entire element

The isotope of an element can determine whether the atom is stable or not 

• Too many neutrons: too much distance → strong nuclear force is weaker

• Too little neutrons: protons too close together → electrostatic force repels

• Need to find a good balance for the atom to be stable 

Isotope Symbol: 

• Top right corner is left blank if there is no charge

• Mass number = Isotope number

• Neutron number = Mass number - Atomic number

Ions

Atoms that have a charge (unbalanced protons & electrons)

• Cations — positively charged ions (lose electrons to make cations)

• Anions — negatively charged ions (gain electrons to make anions)

• Charge = number of protons - number of electrons

• # of electrons defines the ion 

Example Analysis:

Isotope Symbol: 1939K1+

• Protons: 19

• Electrons: 18 (charge is 1+)

• Neutrons: 20 (mass no. - atomic no.)

• Atomic #: 19

• Mass #: 39

• Cation (positive)

Atomic Mass

The weighted average of all naturally occurring isotopes of an element based on their abundance

Calculating Atomic Mass

(Mass of Isotope 1 × % Abundance of Isotope 1) + (Mass of Isotope 2 × % Abundance of Isotope 2) + …

• Mass of Isotope = Isotope Number  = Mass Number

Example Calculation

Isotope 1: Mass = 10.0 amu | Abundance = 70.0% (0.70)

Isotope 2: Mass = 11.0 amu | Abundance = 30.0% (0.30)

Atomic Mass = (10.00.70) + (11.00.30)=10.3

Objectives: 

The Periodic Table (part of Unit 7)

• Identify regions of the periodic table 

  • Use regions to locate elemental information (ex. metals, non-metals, metalloids)

• Metals v.s. Nonmetals (different & similar properties) 

• Understand: metalloids (intermediate properties)

• Orbitals & Principal Energy Levels (relation with regions of the periodic table)

• Define: valence & inner electrons

• Predict the number of valence electrons for elements in different periodic table groups 

Electron Configurations & Orbital Diagrams

• Write electron configurations for elements (& ions)

  • Full & Noble Gas Notation up to the element Carbon 

• Understand Aufbau principle, Hund's rule, and Pauli Exclusion Principle

  • Apply to orbital diagrams 

• Draw orbital diagrams (representing distribution of electrons within orbitals)

Periodic Trends

• Define & Explain: Zeff (effective nuclear charge), atomic radius, ionic radius, ionization energy, electronegativity

• Understand: periodic trends dealing with the properties above

  • How Zeff changes across a period (L→R) & down a group (T→B)

  • Trends in atomic radius, ionization energy, electronegativity using Coulomb’s Law across a period (L→R) & down a group (T→B)

  • Relate periodic trends to the Zeff experienced by valence electrons

• Compare & contrast the sizes of cations (+) & anions (-) with their parent atoms

Electron Excitation (Atomic Emission Spectra Lab)

The Periodic Table

Regions of the Periodic Table

Columns—groups | Rows—periods

• Left of staircase: metals (excluding hydrogen)

• Right of staircase: nonmetals (most are gases)

• Staircase: metalloids (intermediate between metals & nonmetals)

• Hydrogen—nonmetal; in the metal section of the table

• Helium—noble gas (nonmetal); sometimes placed in metal section

• Noble gases: most stable elements (valence shell is full)

• Excluding noble gases, the closer an element is to the sides (L/R) of the periodic table, the more reactive it is 

Metals vs. Nonmetals vs. Metalloids

Metalloids

• Usually look like metals but behave like nonmetals

• Shiny, brittle solids at room temp.

• Semi-conductive 

• Most do not react vigorously with acids 

• Can form alloys with metals

Orbitals & Principal Energy Levels (Shells)

Shells/Principal Energy Levels

Energy levels or electron clouds surrounding the atomic nucleus, where electrons could be found. Each shell has sublevels/orbitals. 

Sublevels/Orbitals

A 3-D area around the atom’s nucleus describing the most likely location of an electron within a given shell. The bigger the shell (or higher the energy level), the more possible orbitals there are. 

• Every orbital can hold up to 2 electrons 

• The 2 electrons must be opposite in spin 

  • Not necessarily directions up & down, but the electrons just need to be opposite in spin

• Each new shell introduces a new, more complex orbital 

  • Shell 1: s orbital

  • Shell 2: p orbital

  • Shell 3: d orbital

  • Shell 4: f orbital

Atomic Orbitals 

• Electrons always want to move with the lowest energy, which means: 

  • Moving in simplest orbital

  • Staying close to the nucleus 

  • Orbitals that require the least energy will always fill up first

    • Sometimes, a simple orbital further away from the nucleus will require less energy than a complex orbital closer to the nucleus

s Orbital

• Sphere shape

  • Simplest; requires least energy (fills up first) 

• 1 configuration → holds a total of 2 electrons 

• Introduced in the 1st shell/principal energy level

  • The only orbital in 1st energy level 

p Orbital

• Peanut shape

  • More complex than s orbital 

  • Shape helps electrons avoid bumping into each other

• 3 configurations → holds a total of 6 electrons (2 electrons per orbital)

• Introduced in the 2nd energy level 

  • 2nd level includes s orbitals & p orbitals

    • s orbital always fills up first 

  • Total of 2+6=8 electrons in 2nd level

• Order of orbitals filled up by electrons from shells 1-2:

  • 1s, 2s, 2p

d Orbital

• Double peanut shape

  • More complex than s & p orbitals 

• 5 configurations → holds a total of 10 electrons 

• Introduced in the 3rd energy level, but electrons do not fill up d-orbital until they have filled up the s-orbital in the 4th energy level

  • Despite increased distance, 4s orbital requires less energy than 3d because of how complicated the d-orbital is

• Order of orbitals filled up by electrons from shells 1-4:

  • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p

f Orbitals

• Flower shape 

  • Most complex orbital 

• 7 configurations → holds a total of 14 electrons

• Appears in the 4th energy level, but electrons do not fill up f orbital until they have filled up the s-orbitals in the 5th & 6th levels & p-orbital in the 5th level

  • Similar idea as d-orbital

  • s & p orbitals further away from nucleus still require less energy than f-orbital close to nucleus because its orbital is so complex

• Order of orbitals filled up by electrons from shells 1-6

  • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p

Orbital Relation with Regions of the Periodic Table

• Orbital blocks indicate orbitals filled by electrons in various elements’ atoms

• Letters (s, p, d, f) represent the last orbital filled up by electrons 

• Numbers (1-7) represent the principal energy level of the last orbital filled

• Columns indicate no. of electrons in the outermost orbital 

• E.g K (4s) → electrons occupy orbitals 1s, 2s, 2p, 3s, 3p, 4s with 1 electron in the 4s orbital | Si (3p) → electrons occupy 1s, 2s, 2p, 3s, 3p with 2 electrons in the 3p orbital

Valence & Inner Electrons

Valence Electrons

Electrons in the outermost shell, or energy level, of an atom. 

• Only found in s & p-orbitals because these orbitals are always at the highest principal energy levels/shells 

Valence electrons influence how an element reacts

• Elements with the same number of valence electrons (same column on periodic table) tend to behave similarly 

• Noble gases are stable because their valence shell is full (unlikely to gain/lose electrons)

Inner Electrons

Electrons in the inner shells of an atom; all non-valence (closest to nucleus)

Stabilize the atom's structure; do not participate in chemical bonding

On the Periodic Table

Electron Configurations & Orbital Diagrams 

Electron Configurations 

Full Electron Configurations: 

A list of numbers and letters describing how electrons are distributed in atoms’ atomic orbitals. 

Leading #: orbital shell (principal energy level)

Letter (s, p, d, f): orbital type 

Superscripts: # of electrons in that specific orbital 

• The sum of superscripts is the total # of electrons of the atom

Examples: Calcium — 1s22s22p63s23p64s2

Boron — 1s22s22p1

Noble Gas (Shorthand) Electron Configurations:

Shorter way of writing electron configurations using the noble gas in the preceding row (period). 

Leading bracket: element symbol of the preceding noble gas

Other characters: electron orbitals on top of the noble gas’s configuration 

Examples: Iron — [Ar] 4s23d6

Silicon — [Ne] 3s23p2

With ions, simply add/subtract total electrons in the configuration to match the charge. 

Example: 

Sodium +1 — 1s22s22p6 or [He]2s22p6

Lewis Dot Diagrams

A way to represent the # of valence electrons in an atom. Dots are drawn around the element symbol to signify electrons. 

• Way #1: start from the top & go around clockwise (to the right)

• Way #2: see it as S-P-P-P structure 

  • Fill up s-orbitals first, rather than going around like in Way #1

Example: Chlorine — 7 valence electrons 

With ions, add/subtract electrons accordingly. Place the diagram in brackets with the charge as a superscript. 

Example: Chlorine -1

Aufbau Principle, Hund’s Rule, and Pauli Exclusion Principle

Aufbau Principle Electrons fill from lowest to highest energy

Pauli Exclusion Principle Each orbital can hold 2 electrons

The 2 electrons have to spin in opposite ways

Hund’s Rule Electrons won’t pair before filling up all the orbitals

Unpaired electrons in the same orbital spin in the same way

Example Application to Orbital Diagrams: 

Incorrect—violates Aufbau Principle

Incorrect—violates Hund’s Rule

Orbital Diagrams 

Orbital diagrams represent the distribution and energy of electrons within the sublevels or orbitals of an atom. Apply the three rules above. 

Example: 

Nitrogen (ground-state) Electron Configuration — 7 electrons

Periodic Trends

Zeff, Atomic Radius (Ionic Radius), Ionization Energy, Electronegativity 

Zeff (effective nuclear charge): 

Charge of the nucleus felt by its valence electrons 

Calculation: #protons-#inner electrons

• Inner electrons: all non-valence electrons 

  • A.k.a 'shielding electrons’ → shields valence electrons from the full force of the nucleus 

  • For ICPS, we consider all inner electrons to be equal 

  • Example: 1s22s22p6 — shell 2 is outermost; only 2 inner electrons

• No units for Zeff

Examples:

1. Sodium (Na) 

• #protons: 11 #electrons: 11 (for neutral atoms)

• #inner electrons: 10

Zeff=11-10=1

• 1 is also #valence electrons 

• For most uncharged atoms, Zeff = # valence electrons

2. Sodium +1 (Na+1)

• #protons: 11 #electrons: 10

• #inner electrons: 2

Zeff=11-2=9

Atomic Radius: 

Radius of electron cloud (distance from the centre of the nucleus to edge of electron cloud)

Ionic Radius: radius of electron cloud of an ion 

• Differs from its parent atom because # of electrons changes, causing the distance of the shell from the nucleus to differ as well

Ionisation Energy:

Energy needed to break the electron free from the force felt by nucleus

• Stronger Fe (electrostatic force) = greater ionisation energy

• Changes depending on which electron you’re trying to take off

Electronegativity: 

The attractive force one atom’s nucleus has for another atom’s valence electrons when the two atoms are covalently bonding 

• Only comes into play when atoms are covalently bonding 

• Typically doesn’t deal with noble gases because they don’t like to bond

Periodic Trends (Zeff-related)

Zeff Trends: 

L → R (across a period)

• Increases → #p increases while #inner e stays the same 

  • Same outermost shell, so # inner electrons don’t change 

T → B (down a group)

• Stays the same → #p and #inner e increase at the same rate 

Atomic Radius

Fe=kQ1Q2r2 (Coulomb’s Law applied to atoms)

     Q1—valence electrons | Q2—Zeff (nucleus) | r—radius using shells (not actual radius)

L → R (across a period)

• Decreases → Fe=k(same)(up)(same)  → Fe increases; AR decreases

  • Charge of valence electron always stays; Zeff increases (#p increases); outermost shell stays (same distance)

  • Radius decreases as force increases (pulls tighter)

T → B (down a group)

• Increases → Fe=k(same)(same)(up) → Fe decreases; AR increases

  • Zeff stays (shown above); outermost shell increases (moves up a principal energy level each period) 

  • Force decreases, so radius increases 

Ionisation Energy

Fe=kQ1Q2r2

Q1–valence electrons | Q2–Zeff (nucleus) | r–atomic radius

L → R (across a period)

• Increases → Fe=k(same)(up)(down)  → Fe increases; IE increases

  • Zeff increases (#p increases); atomic radius decreases 

  • Force increases; ionisation energy increases (tighter hold)

T → B (down a group)

• Decreases → Fe=k(same)(same)(up) → Fe decreases; IE decreases

  • Zeff stays; atomic radius increases

  • Force decreases; ionisation energy decreases

Electronegativity 

Fe=kQ1Q2r2

Q1–valence electrons of 2nd atom | Q2–Zeff (nucleus of 1st atom on 2nd atom’s electrons) | r–atomic radius of 1st atom

L → R (across a period)

• Increases → Fe=k(same)(up)(down)  → Fe increases; EN increases

  • Zeff increases (#p increases); atomic radius decreases 

  • Force increases; electronegativity increases (attraction between two atoms)

T → B (down a group)

• Decreases → Fe=k(same)(same)(up) → Fe decreases; EN decreases

  • Zeff stays; atomic radius increases

  • Force decreases; electronegativity also decreases

• Not a periodic trend → only exists when two atoms are covalently bonding

Cation & Anion Sizes

Cations (positive; lose electrons)

• Radii tend to get significantly smaller 

  • Atoms likely to lose electrons tend to lose a whole shell

  • Shells held much tighter (high Zeff)

Anions (negative; gain electrons)

• Radii tends to get slightly bigger 

  • Atoms likely to gain electrons tend to have electrons added to the same shell

  • Electrons repel each other slightly more, pushing the radii out, but Zeff stays the same 

Electron Excitation (Lab)

Ground State: electrons at the lowest possible energy level

Excited State: electrons at a higher-energy shell 

• Happens by absorbing a photon, heat, or electricity 

Relaxing: electrons exiting excited states by losing energy, emitting a photon of light

• Photons emitted are determined by spacing between excited & ground state

  • Various photons are different wavelengths of light, determining color

• Greater spacing = shorter wavelength (violet side of color spectrum)

• Wavelengths/color emitted differ for each element 

  • Number of electrons & spacing between shells differ for each element

Spectroscope: reveals atomic emission spectrum, consisting of lines of color (emitted colors)

• Unique for every element

 Objectives: 

Covalent Bonding

• Define: Covalent Bonding 

• Naming covalent compounds (using prefixes)

• Single, double, and triple bonds 

  • Relate bond order to bond strength to bond length 

• Draw Lewis structures with resonance structures 

• Molecule shapes & geometry using VSEPR 

• 3-D representations of molecules using VSEPR 

• Polar v.s. Nonpolar covalent bonds (based on electronegativity)

  • Relate molecular shape to molecular polarity

Intermolecular Forces 

• Describe Intermolecular Forces (LDF, dipole-dipole, hydrogen bonds)

• Relation to physical properties (boiling/melting points, solubility, etc) 

• IMF Lab

Covalent Bonding

Covalent Bonding 

The mutual attraction between 2 atoms’ nuclei for each other’s valence electrons. 

• Only happens between nonmetals 

• No. of covalent bonds an atom forms depends on the no. of electrons needed to fill its valence shell

  • Hydrogen & Halogens only form one covalent bond (only needs 1 more electron to fill their valence shell)

  • Bonding only involves valence electrons

Diagram: 

Explanation: 

As atoms approach one another, their electrons start detecting each other, creating slight repulsive forces (too slight to matter). As the atoms get closer, their positive nuclei attract each other’s electrons. This attractive force decreases potential energy, as shown in the diagram. The bottom of the curve indicates the distance at which the bonded atoms are most stable (lowest energy). The magnitude of the potential energy at that point is how much energy is needed to break the covalent bond. If the atoms move closer, their positive nuclei start to repel. This strong repulsive force dominates the attractive forces, causing a spike in potential energy. Eventually, the repulsive force will break the covalent bond. 

• Negative potential energy indicates that energy must be provided to break the bond

• Positive potential energy indicates that energy must be provided to form the bond

Naming Covalent Compounds

Only binary covalent compounds (2 types of atoms)

1. 1st element: regular name (e.g Cl→chlorine; O→oxygen)

2. 2nd element: change to ‘-ide’ name (e.g Cl→chloride; O→oxide)

3. Add prefixes to indicate number of atoms 

   1. *don’t use ‘mono’ before 1st element

All -ide names

• Only nonmetals (excluding noble gases) covalently bond:

Hydrogen — Hydride | Carbon — Carbide | Nitrogen — Nitride | Oxygen — Oxide

Fluorine — Fluoride | Phosphorus — Phosphide | Sulfur — Sulfide 

Chlorine — Chloride | Selenium — Selenide | Bromine — Bromide | Iodine — Iodide

Prefixes

Examples

N2O Dinitrogen Monoxide

CO2 Carbon Dioxide

Carbon tetrachloride CCl4

Sulfur trioxide SO3

Single, Double, and Triple Bonds

Single Bonds

• One shared pair of electrons (two electrons)

  • (X-X) 

• Lowest energy/strength → longer length

Double Bond 

• Two shared pairs of electrons (four electrons)

  • (X=X) 

Triple Bond 

• Three shared pairs of electrons (six electrons)

  • (X≡X) 

• Highest energy/strength → shorter length 

Bond Order: number of bonds between a pair of atoms (single, double, triple)

Bond Length: average distance between nuclei of two bonded atoms 

• An average because the atoms are constantly moving & changing distance

• X-coordinate of the bottom of the curve in the diagram above

Bond Energy/Strength: average amount of energy required to break the bond

• An average because it varies depending on the specific molecule the bond is in

Note: can only compare bond order within the same elements (can’t compare hydrogen triple-bond to nitrogen single-bond because the element changes bond length & strength)

Lewis Structures 

• 2-D representation of the bonding between atoms (covalent compounds only)

  • Only deals with valence electrons (includes bonding pairs & lone pairs)

How to Draw:

1. Count valence electrons in the molecule

2. Identify central atom

   1. Least electronegative atom (least attractive to electrons)

   2. Usually the first atom written 

   3. Never hydrogen or halogens (these can only form 1 covalent bond because their valence shell is 1 away from being full)

3. Create a skeletal structure 

   1. Use lines to indicate bonds; each line represents 1 electron pair

      1. Double bonds have 2 lines; triple bonds have 3 lines

   2. Four atoms around  central atom 

4. Distribute remaining electrons 

   1. Add lone pairs or create double/triple bonds 

   2. Each atom needs a full valence shell (eight electrons, except hydrogen)

      1. Each bond (one line) is 2 electrons 

   3. Total # of electrons on diagram needs to equal total valence electrons

Examples (simple molecules):

H2O

1. Valence electrons: 8 

   1. 2 hydrogen & 1 oxygen

   2. H: 1 valence O: 6 valence 

   3. Add together: 21+16=2+6=8

2. Central Atom: Oxygen (hydrogen cannot be central atom)

3. Skeletal Structure:

4. Remaining Electrons:

SCl2

1. Valence electrons: 20

1. 1 Sulfur & 2 Chlorine

2. S: 6 valence Cl: 7 valence 

3. Add together: 16+27=6+14=20

5. Central Atom: Sulfur (less electronegative)

6. Skeletal Structure:

7. Remaining Electrons:

Lewis Structures for Polyatomic Ions 

Polyatomic Ions: covalently bonded molecules with a charge

• Add brackets around the diagram & write charge on top right corner

Example: 

(element: CO3 | charge: 2-)

Resonance

When there’s a double or triple bond that can be in multiple locations. The actual structure of a molecule with resonance is a hybrid of all the resonance structures. 

Indicating Resonance

• Draw all resonance structures 

• Connect the diagrams with double arrows 

Example:

Element: NO3- Resonance: 3

Reasonable Resonance Structures (Lab): 

Resonance structures are equally likely to occur (electrons constantly change positions between the possible structures). 

Reasonable Example:

SO₂ 

• Both oxygens have the same electronegativity (same pull on electrons)

  • One oxygen does not attract electrons more than the other

  • Electrons equally likely to be at either oxygen

• Both resonance structures equally likely to occur

• Bond length: one and a half (electrons constantly switching positions, so bond length is the average)

Unreasonable Example:

CO₂  

• Structures 2 & 3 are unlikely to occur

  • Both oxygens have equal electronegativity — electrons are unlikely to gather near one oxygen 

• Structure 1 is the most likely → unreasonable resonance structures

  • Electrons likely to spread out equally between two oxygens

VSEPR (Valence Shell Electron Pair Repulsion Theory)

• Theory telling us how electrons repel during bonding

• Used to determine the 3-D geometry & shape of molecular structures

  • Lewis-structure is only 2-D

Theory: 

1. Electron groups (lone pairs & bonding pairs) repel (negative on negative) and will arrange themselves in a way that maximises distance between them. 

• Creates biggest angle possible between electron groups

2. Lone pairs repel more than bonding pairs

• Electrons in lone pairs repel each other more

• Lone pairs repel bonding pairs more

3. Angle of molecular structure varies based on lone pairs

• With lone pairs, the angle between bonding pairs are less than the base angle because lone pairs repel more and take up a greater angle, forcing the bonding pairs closer together (smaller angle)

Lewis v.s. VSEPR

Molecular Geometry

• Examines the location of electrons on the central atom(s)

Steps:

1. Draw Lewis Structure

2. Determine number of electron locations (on central atom)

   1. One bond location (single, double, or triple) is one location

   2. One lone pair is one location

3. Name the geometry depending on the number of electron locations

Types: 

Molecular Shapes

• Examines electron locations relative to each other (#bonding locations v.s. #lone pairs)

Steps:

1. Draw Lewis Structure

2. Determine number of bonding locations v.s. lone pairs

3. Name the shape

Types: 

VSEPR Relation

• VSEPR theory describes the 3-D shape 

  • Linear: no lone pair; bonding pairs repel to create biggest angle (180º); a straight line

  • Bent: lone pair pushes away two bonding pairs, creating a bend

  • Trigonal planar: no lone pair; bonding pairs repel to create biggest angle (120º); exists on one plane

  • Trigonal pyramidal: lone pair pushes away three bonding pairs, creating pyramid-like shape

  • Tetrahedral: no lone pair; bonding pairs repel to create biggest angle (109.5º)

3-D Representations 

• Depict the molecule’s 3-D shape 

  • Lone pairs are NOT drawn (only bonds)

  • Keeps as many connected atoms on one plane as possible

• Wedges

  • Indicates the bonded atom is on a different plane

  • Different colored wedges: atoms on different planes (only tetrahedrals)

Polarity

Polar Covalent Bonds: 

When 2 bonded atoms have differing electronegativities (one atom pulls on electrons harder than the other)

• Electrons spend more time closer to the more electronegative atom

  • Creates partial charge 

  • More electronegative atom: negative partial charge 

  • Less electronegative atom: positive partial charge

• For ICPS: bonded atoms with an electronegativity difference greater than 0.35 create polar bonds 

Dipole Depiction

Nonpolar Covalent Bonds:

When the difference in the electronegativity of 2 bonded atoms is less than or equal to 0.35

• CH (carbon–hydrogen) bonds are not polar 

  • C: 2.55 | O: 2.20 (difference in EN = 0.35)

Molecular Polarity

A molecule is polar if there is an overall partial charge on one end of the molecule (when the dipoles don’t cancel out)

• Depends on molecular shape

Examples

Polar:

Nonpolar: (all CH bonds, which are nonpolar)

Intermolecular Forces

Intermolecular Forces: Forces between molecules

Strength of IMFs LDF (weakest) → Dipole-Dipole → Hydrogen Bond (strongest)

London Dispersion Forces (LDF)

Temporary attractive force between opposite temporary dipoles of molecules. LDF is a relatively weak intermolecular force. 

Temporary Dipoles

• When a molecule’s electrons are dispersed unevenly, causing more of them to gather on one side, creating temporary partial charges/dipoles

  • More electrons=negative partial charge

• Opposite ends of temporary dipoles on two molecules can attract

• More electrons = Greater LDF 

  • The more electrons a molecule has, the more unevenly dispersed they are likely to be, creating greater temporary partial charges

Dipole-Dipole Interactions

Attractive force between opposite dipoles of molecules (positive & negative end); force is stronger & more permanent than LDF

• Molecules need to be polar (needs permanent dipoles that do not cancel out)

Hydrogen Bonding

A type of dipole-dipole interaction that is particularly strong (not actually a bond, but an intermolecular force) 

Molecules that hydrogen bond have strong partial charges because of larger electronegativity differences between atoms in the molecule, making the attractive intermolecular force stronger. 

Molecules must have N–H, O–H, or F–H bond (hence “hydrogen bond”)

• Hydrogen has a low electronegativity (2.20), allowing for greater EN differences with other atoms it covalently bonds with

• Greater EN → stronger partial charges → stronger dipole-dipole force

Physical Properties 

Intermolecular forces bind molecules together. Their strength determines the state of the compound (solid, liquid, gas). 

• Solids: strong IMF; molecules are held tightly together — requires more energy to overcome IMF & break molecules apart

• Gases: weak IMF; molecules held loosely together — requires little energy to overcome IMF & break molecules apart

Phase changes: changing between solid, liquid, and gas states

• Putting in energy to break IMF can make solids become liquid or gas

• Taking away energy can make gases become liquid or solid

Boiling/Melting points

• Stronger IMF = higher boiling/melting points

  • Boiling & melting points are temperatures where substances become gas or liquid 

  • Stronger IMF means more energy is required to break forces & separate molecules → higher boiling/melting points (more heat energy)

Solubility

• Likes dissolve likes

  • Nonpolar compounds dissolve in nonpolar solutions

  • Polar compounds dissolve in polar solutions

Evaporation Rate (IMF Lab)

• A cooling process—the rate at which a liquid becomes gas and escapes into the air (a phase change)

• Can be measured by cooling rate (∆T/∆t) — change in temp/change in time

  • In the evaporation process, energy is provided to break the IMFs holding the compound in its liquid state, converting it to the higher-energy state of vapor (gas) 

  • As liquid evaporates, the higher-energy gas molecules (high energy=high temp) escape into the air, taking the energy with it

  • Lower-energy molecules (low temp) remain, creating a cooling effect 

• Stronger IMF = slower evaporation rate

  • Requires more energy to break IMF & allow liquids to vaporise (become gas)

IMF Lab

R2 — indicates how well the trendline models the data 

• R2 = 1: line of regression perfectly fits the data

  • Variation in the dependent variable is explained by the independent variables 

• R2 = 0: line of regression does not explain the data at all

  • Variations in the dependent variable are not caused by the independent variables at all

• Low R2 — there are other variables (extraneous variables) apart from the considered independent variables influencing the dependent variable 

• High R2 — the independent variables considered are the main factors influencing the dependent variable

Extraneous Variables

Factors that may influence the dependent variable or controlled conditions but are not considered an independent variable in the experiment. 

Objectives: 

• Define: Ionic bonding

• Formulas of ionic compounds (using ion charges)

• Naming ionic compounds (binary, polyatomic ions, transition metals) 

• Ionic bond strength (related to charge & distance—Coulomb’s Law)

• Properties of ionic compounds (boiling/melting points, solubility) 

Ionic Bonding Definition 

Definition: 

A chemical bond formed by the attraction between a cation (positively charged) and an anion (negatively charged). Ionic bonding happens between metals and nonmetals. 

• Metals tend to lose electrons and form cations. 

• Nonmetals tend to gain electrons and form anions. 

• Polyatomic ions are groups of atoms that collectively act as an ion. 

Formulas of Ionic Compounds

Ionic bonds form lattice structures, which are expressed as formula units. 

Lattices: 3D arrangement of oppositely charged ions held by ionic bonds 

• Every cation is bonded to every anion near it

Formula Units: gives the ratio of various types of atoms in an ionic compound

Writing Formula Units

1. Determine the charges of the ions 

2. Place the cation first 

3. Use criss-cross method 

   1. Charge of the cation becomes the subscript of the anion, and vice versa

2. For polyatomics, include parentheses if writing a subscript

4. Simplify to the lowest multiple 

   1. Example: Mg2O2 becomes MgO

Covalent vs. Ionic

Similarities:

• Forms from electrostatic forces 

Naming Binary Ionic Compounds

1. Metal (cation) goes first

2. Cation’s full name (e.g sodium, aluminium)

3. Followed by anion’s name in -ide form (e.g oxide, fluoride)

Transition Metals 

• Charge will be given (not consistent)

  • If given a formula unit: find charge by seeing which value would make the charges balance 

  • E.g FeCl3 – Iron (transition metal) & chlorine (1- charge); there is 1 iron and 3 chlorine atoms, so the charge of 3 chlorine atoms must equal the charge of 1 iron atom. 1*3=3, so the iron in this compound has a charge of 3+. 

• Comes first, because they are metals (cations)

• Add roman numerals in parentheses to indicate their specific charge 

  • Iron (III) chloride; Lead (II) oxide

Polyatomic Ions

• If cation (positive), comes first

• If anion (negative), comes second

• State full name for all polyatmoic ions (e.g nitrate, phosphate)

Examples

1. Lead (II) sulfite — PbSO3

2. Iron (III) oxide — Fe2O3

3. Re(NO2)3 — Rhenium (III) nitrate

Ionic Bond Strength 

Ionic bonds are held by electrostatic force (attraction between two opposite charges), so Coulomb’s Law applies. More force means stronger bonds (directly related). 

Factors Affecting Ionic Bond Strength:

(involves charge & distance)

Ion Charge: directly related to bond strength

(greater charge, greater force)

Ion Size: inversely related to bond strength

(greater size, greater distance, weaker force)

Lattice Energy

The energy required to break all of the bonds in a lattice (all the ionic bonds)

Bond strength and lattice energy are directly related 

(stronger ionic bonds = higher lattice energy)

Properties of Ionic Bonds 

Melting/Boiling points

Compounds melt or boil when they change from solid to liquid or liquid to gas. This involves breaking bonds that hold the atoms together. 

Stronger bonds = more energy to separate particles = higher melting/boiling points

Covalent vs. Ionic

Covalent compounds: formed by molecules; IMFs need to be overcome to separate molecules

Ionic compounds: lattice structures; ionic bonds need to be broken to separate ions. 

Ionic compounds have higher melting/boiling points

• Ionic bonds are much stronger & harder to break than IMFs, as they deal with full charges while IMFs deal with partial charges (charge)

• Distance between molecules held by IMFs is greater than distance between atoms ionically bonded (distance)

Solubility (Ionic bonding & Solubility lab)

Ionic compounds dissolve in water when the ion-dipole interaction (interaction between charges on ions and dipoles on water molecules) is stronger than their ionic bonds. To compare the two interactions, consider charge & distance (Coulomb’s Law). 

Ion-dipole interaction:

Water molecules are polar. They have partial charges formed by the unequal sharing of electrons between hydrogen and oxygen atoms. 

Oxygen is much more electronegative than hydrogen, so electrons stay closer to oxygen, creating a slight negative charge on the oxygen end of water molecules. 

This partial charge interacts with the charged ions in ionic compounds. 

Stronger ionic bonds = Less soluble

Precipitates: 

Insoluble solids that remain as solids in a liquid solution 

Hydration Sphere:

A layer of water molecules that surrounds an ion when the ionic compound is dissolved in water

Objectives: 

Chemical Reactions

• Understand chemical reactions 

  • Involve breaking & forming of bonds

  • Resulting in the formation of new compounds 

• Law of Conservation of Mass (CoM)

• Write skeletal equations

• Diatomics (HOFBRINCL)

• Balancing chemical equations 

  • Relate to Law of CoM

Stoichiometry 

• Solving stoichiometry problems with mole ratios

• Using mole conversion factors 

  • Molar mass

  • Avogadro’s #

• Limiting & excess reactants 

• Percent yield & percent error 

Chemical Reactions

Chemical Reactions Overview 

• Involves both breaking & forming bonds

• Results in the formation of new compounds

Example: chemical reaction forming water

Terminology:

Reactants: what you start with (left of the arrow)

Products: what you end with (right of the arrow)

Solution/Aqueous: dissolved in water 

Law of Conservation of Mass (CoM)

Matter can never be created nor destroyed 

• Applies to the universe or closed systems

Open system: matter can leave or enter a system 

Closed system: matter cannot leave or enter a system 

• In the real world, closed systems are never perfect

• Percent error & yield account for matter lost/gained in a system (explained below)

Skeletal Equations 

An unbalanced chemical equation showing the formulas of the reactions and products. 

• Includes states of matter (solid, liquid, gas, aqueous)

• Ensures all types of elements on 1 side appear on the other (law of CoM)

  • Equations do not have to be balanced

Format:

Reactant (state)  +  …  + Reactant (state)   →   Product (state) + …  +  Product (state)

(s): solid (l): liquid (g): gas (aq): aqueous (dissolved in water)

Example: CH4 (g) + O2 (g) → CO2 (g) + H2O (l)

Diatomics (HOFBrINCl)

Elements that only exist as molecules composed of two of the same atom (e.g oxygen atoms are only found as O2 molecules, never as a singular O atom). 

1. Hydrogen (H)

2. Oxygen (O)

3. Fluorine (F)

4. Bromine (Br)

5. Iodine (I)

6. Nitrogen (N)

7. Chlorine (Cl) 

Balancing Chemical Equations

A chemical equation is balanced if the number of each type of atom in the reactants equals the number of each type of atom in the products. 

How to Balance Equations:

• Change the coefficients to get equal atoms on both sides (cannot change subscripts)

• Simplify

• In difficult questions, polyatomic ions could break up into separate elements

Examples: 

Unbalanced: CH₄ (g) + O₂ (g)     →     CO₂ (g) + H₂O (g)

(1C, 4H, 2O     →     1C, 3O, 2H)

Balanced: CH₄ (g) + 2O₂ (g)     →     CO₂ (g) + 2H₂O (g)

(IC, 4H, 4O     →     1C, 4O, 4H)

Unbalanced: H₂ (g) + O₂ (g)    →     H₂O (l) (2H, 2O → 2H, 1O)

Balanced: 2H₂ (g) + O₂ (g)     →     2H₂O (l) (4H, 2O → 4H, 2O)

Relation to Law of CoM

Atoms (matter) cannot be created or destroyed, so atoms cannot be gained or lost in a chemical equation representing a reaction. By balancing equations, we ensure that the reaction agrees with the Law of CoM. 

Stoichiometry

Stoichiometry: Relationship between quantities of reactants & products in a chemical reaction, typically involving a ratio of integers (switching between substances essentially). 

Moles & Molar Mass

Mole: 6.021023 particles of something (an element, molecule, etc)

• A counting number (standard unit for measuring the amount of a substance)

• 6.021023 is Avogadro’s Number

Conversion Factor: 

1 mole (of a substance)   =   6.021023 particles (of a substance)

Molar Mass: amount of grams in 1 mole of a substance (units: grams/mole)

• Indicated by atomic mass on the periodic table (number below the element)

  • Atomic mass is the weighted average of the masses of all of an element’s naturally occurring isotopes (expressed in amu)

Example:

Molar Mass of Sulfur  =  32.066 grams/mole

• 32.066 grams of sulfur atoms are in 1 mole of sulfur

• Molar mass of molecules: the sum of the molar masses of each involved atom

Example: H₂O molar mass

Atoms 2H, 1O

Molar Masses H—1.008 g/mol | O—16.00 g/mol

Calculation: 21.008+16.00=18.02 g/mol

Conversion Factor: 

1 mole (of a substance)   =    Indicated Mass (g) (of a substance)

Mole Ratios 

The ratio between the amount in moles of two substances involved in a balanced chemical reaction. 

• Always BALANCE equations before determining mole ratio

What mole ratio tells us:

For a specific chemical reaction, if you have a certain amount of 1 substance, how much of another substance will be created (if it is a product) or used up (if it is another reactant) through the reaction, assuming that there is the perfect amount of other substances to make the reaction work. 

Example

Balanced Equation: 2H₂ + O₂     →     2H₂O

Mole Ratio of:

H₂ and H₂O          2 mol H22 mol H2O     every 2 moles of H₂ makes 2 moles of H₂O

O₂ and H₂O          1 mol O22 mol H2O         every 1 mole of O₂ makes 2 moles of H₂O

H₂ and O₂             2 mol H21 mol O2           every 2 moles of H₂ used up, 1 mole of O₂ is used up

Mole Conversions (all cases)

• Dimensional Analysis

Using Molar Mass (g/mol): 

Grams → Moles ___ grams*1 mol ___  g=___ moles

Moles → Grams ___ moles*___  g 1 mol=___ grams

Using Avogadro’s Number: 

Moles → Particles ___ moles* 6.021023 particles1 mol=___ particles

Particles → Moles ___ particles* 1 mol6.021023 particles=___ moles

Using Mole Ratio:

Mole (substance 1) → Mole (substance 2)

___  moles (substance 1)*___ mol2___ mol1=___  moles (substance 2)

Stoichiometry Problems (Examples)

• Deals with going between various substances

1-Step Conversion (Mole:Mole)

• Uses mole ratio only

  • Goes from moles of one substance to another

Example: 

How many moles of hydrogen are required to produce 2.45 moles of ammonia?

Skeletal Equation: N₂ + H₂     →     NH₃

1. Balance Equation: N₂ + 3H₂     →     2NH₃

2. Determine mole ratio: 3 mol H22 mol NH3

3. Dimensional Analysis: 

2.45 mol NH33 mol H22 mol NH3=3.68 mol H2

4. Answer: 3.68 moles of hydrogen are required. 

2-Step Conversion (Mole:Mass)

• Uses mole ratio & molar mass

  • Goes from mass to mole of one substance, then moles of that substance to moles of a second substance 

  • Order can be switched around based on question 

Example: 

How many moles of zinc chloride are formed when 3.61g of hydrochloric acid reacts?

Skeletal Equation: Zn + HCl     →     ZnCl₂ + H₂

1. Balance Equation: Zn + 2HCl     →     ZnCl₂ + H₂

2. Determine molar mass of hydrochloric acid (HCl):

1.008 g/mol H + 35.453 g/mol Cl   =   36.461 g/mol HCl

3. Determine mole ratio: 1 mol ZnCl22 mol HCl

4. Dimensional Analysis:

3.61g HCl1 mol HCl36.461g HCl1 mol ZnCl22 mol HCl=0.0495 mol ZnCl2

5. Answer: 0.0495 moles of zinc chloride are formed. 

3-Step Conversion (Mass:Mass)

• Uses mole ratio & molar mass

  • Goes from mass to mole of one substance, then moles of that substance to moles of a second substance, then moles to mass of the second substance 

Example: 

How many grams of ammonia are produced if 9.50g of nitrogen are available? 

Skeletal Equation: N₂ + H₂     →     NH₃

1. Balance Equation: N₂ + 3H₂     →     2NH₃

2. Determine molar masses:

N₂     —     214.01=28.02 g/mol N2

NH₃     —     14.01+31.008=17.03 g/mol NH3

3. Determine mole ratio: 2 mol NH31 mol N2

4. Dimensional Analysis:

9.50g N21 mol N228.02g N22 mol NH31 mol N217.03g NH31 mol NH3=11.6g NH3

5. Answer: 11.5 grams of ammonia are produced. 

Limiting & Excess Reactants

Limiting Reactant: the reactant that will be completely used up

Excess Reactant: the reactant that will have leftovers after the reaction

Determining the Limiting Reactant:

1. Balance the Equation

2. Find mass of product formed by each reactant based on the reactants’ masses

   1. Use dimensional analysis to work from the mass of one substance to another

3. The reactant that produces the least amount of product is the limiting reactant 

   1. The other is the excess reactant 

Determining the Amount of “Excess” Reactant Leftover After the Reaction: 

1. Find difference between the mass of product produced by each reactant 

   1. Difference is the amount of excess product that excess reactant can produce 

2. Take excess mass of product; convert to mass of excess reactant 

   1. Use dimensional analysis to work from mass to mass 

3. The mass is the amount of leftover excess reactant after the reaction

Example

Copper (II) chloride reacts with sodium nitrate to form copper nitrate and sodium chloride. If 15.0g of copper (II) chloride reacts with 20.0g of sodium nitrate, which is the limiting reactant? How much sodium chloride can be formed?

1. Chemical Equation

Skeletal: CuCl₂ + NaNO₃     →     Cu(NO₃)₂ + NaCl

Balanced: CuCl₂ + 2 NaNO₃     →     Cu(NO₃)₂ + 2 NaCl

2. Find mass of product produced by each reactant based on given masses

   1. Use dimensional analysis (3-step conversion stoichiometry problem)

      1. Find Molar Masses: 

CuCl₂ — 134.45g/mol CuCl₂

NaNO₃ — 85.0g/mol NaNO₃

NaCl — 58.44g/mol NaCl

2. Find Mole Ratio: 2 mol NaCl1 mol CuCl2 2 mol NaCl2 mol NaNO3

3. Dimensional Analysis

Copper (II) chloride:

15.0g CuCl21 mol CuCl2134.45g CuCl22 mol NaCl1 mol CuCl258.44g NaCl1 mol NaCl=13.0g NaCl

Sodium nitrate:

20.0g NaNO31 mol NaNO385.0g NaNO32 mol NaCl2 mol NaNO358.44g NaCl1 mol NaCl=13.8g NaCl

3. Answer: CuCl₂ is the limiting reactant (can produce less product); 13.0 grams of sodium chloride can be formed. 

Percent Yield & Percent Error 

Percent Error

|Experimental Value - Theoretical  Value|Theoretical Value100

Percent Yield

Experimental YieldTheoretical Yield100

• >100%    —    gained matter

• <100%    —    lost matter

Relationship to Law of CoM

If matter should theoretically be conserved within a system, percent error and percent yield show how much experimental results agree with the Law of CoM (low error, high yield is generally desired). 

Copper Cycle Lab

• Applying Law of CoM (Copper conserved in the system, but changes in phase)

• Stoichiometry calculations 

• Balancing & Analysing Chemical Equations

  • Precipitates, physical observations (e.g color, steam)

Objectives: 

Thermal Energy

• Understand: heat, thermal energy, and temperature

• Different ways of heat transfer

• Define: enthalpy 

Enthalpy of Physical Changes 

• Understand: solids, liquids, and gases (macroscopic & microscopic level)

• Calculate: energy changes in heating & cooling curves

Enthalpy of Chemical Reactions

• Endothermic v.s Exothermic reactions 

  • Represent using energy diagrams

• Define: activation energy 

• Describe: collision theory (collision, orientation, sufficient KE)

• Explain: relationship between bond breaking & forming determines overall enthalpy change 

• Thermochemical equations (includes balanced equation & enthalpy change)

• Determine: enthalpy change of a reaction by designing & analysing calorimetry experiments 

Mathematical Calculations 

• Mathematically describe heat flow in a calorimetry experiment 

• Calculate: heat absorbed/released for a given amount of reactant or product

Thermal Energy

Heat, Thermal Energy, and Temperature 

Thermal Energy

Sum of the kinetic energy (KE) of all particles in a substance. 

• KE comes from the movement of particles

• Wording: A system has thermal energy. 

Heat 

Amount of thermal energy transferred in or out of a system

• Wording: heat energy is removed/added to a system. 

Heat always flows from a place with more thermal energy to less (hot to cold) until a thermal equilibrium is reached (equal temperature). 

• At thermal equilibrium, there is still heat flow, but it is equal in both directions → no net flow of heat

Temperature 

Average kinetic energy of the particles in a substance 

KE=32kT T=temperature | k=constant 

• Kinetic energy is directly proportional to temperature (more particle movement, higher temperature)

Temperature Scales

1. Celsius ºC=(ºF-32)59

2. Fahrenheit ºF=(ºC95)+32

*change in 1ºF change in 1ºC

3. Kelvin K= ºC+273.15

Heat Transfer

Conduction: heat transfer through direct contact (when molecules touch)

• Most often through solids

• Molecules exchange kinetic energy by bumping into each other

• Heat flows from 1 matter to another 

Convection: heat transfer through the movement of matter (currents)

• Only happens in fluids (gas or liquid)

• Heat moves with the matter as it flows 

Radiation: heat transfer through electromagnetic waves 

• Heat does not have to pass through matter (can transfer through empty space)

• Heat is sent through waves 

Enthalpy

Heat energy absorbed or released by a system during a chemical or physical process

• Know: Enthalpy related to temperature change & Enthalpy related to phase change (explained more later)

Calorimetry 

Process of measuring the heat energy released or absorbed in a system. 

Assumption in calorimetry experiments: No energy is lost to the surroundings 

Equation: Q1=-Q2

(energy gained in system = energy lost in surroundings)

Enthalpy of Physical Changes

Solids, Liquids, and Gases (including important diagrams)

Phase Changes

• Fusion: between solid & liquid (melting/freezing)

• Vaporisation: between liquid & gas

Phase Diagrams 

• Triple point: specific temperature & pressure where all three states of a substance can coexist at equilibrium

• Phase diagram shows how pressure (atm) & temperature (ºC) influence phase changes

Enthalpy related to Physical Changes

Enthalpy related to Phase Change

Q=Hm

• Q — energy change (J)

• H — enthalpy of phase change (J/g)     (Hfus or Hvap)

  • Amount of energy absorbed or released during a phase change 

  • Differs for every compound and its every phase

• m — mass (g)

Enthalpy related to Temperature Change

Q=mcT

• Q — energy change (J)

• m — mass (g)

• c — specific heat capacity (J/gºC)

  • Energy required to raise the temperature of 1g of a substance by 1°C

  • Differs for every compound and its every phase

• T — change in temperature (ºC)

Heating & Cooling Curves

• Heating Curve: Energy is put into the system over time, changing either the substance’s temperature or phase (graph above going from left to right)

• Cooling Curve: Energy is being released from the system over time (graph above going from right to left)

• The system absorbs heat, which becomes either KE/PE

  • A→B, C→D, E→F: changes in temperature → increasing KE as it is a measure of temperature (KE=32kT)

  • B→C, D→E: changes in phase → the distance between particles increases, causing an increase in the system’s PE

• The enthalpy calculation for each section on the diagram is given by the equation for Q below the x-axis

Enthalpy of Chemical Reactions

Exothermic v.s Endothermic Reactions

Exothermic Reaction     —     releases heat into the surroundings

• Enthalpy change (Hrxn) is negative (-)

Endothermic Reaction     —     absorbs heat from the surroundings 

• Enthalpy change (Hrxn) is positive (+)

Equation

Hrxn=Hproducts-Hreactants

Reaction Energy Diagrams

Hrxn > 0 Hrxn < 0

• Transition State (Activated Complex): the substance’s state when it has the amount of energy indicated by the peak of the graph (highest energy point)

  • Required state for reactants to start reacting (breaking/forming bonds)

• Enthalpy of Transition State (H*): energy of the transition state (y-value at the peak)

• Specific Reactant/Product Labels: include A+B (reactants), C+D (products)

Activation Energy 

The minimum amount of energy required to make reactants start interacting. 

• Activation energy (Ea) is not related to the enthalpy change (H) 

Equation: Ea=H*-Hreactants (H* = Enthalpy of Transition State)

Collision Theory

A model explaining how chemical reactions occur at the molecular level. 

Three conditions must be met for a collision to be successful: 

1)     Collision Molecules must collide & come in contact with each other

2)     Correct Orientation Molecules must collide in the right orientation

3)     Sufficient KE Molecules must have enough energy to reach activation energy

Bond Breaking & Bond Formation

Bond Breaking

• Endothermic — energy put in to break the bond; reaction absorbs energy

  • Particles separate, increasing PE (energy needs to be put in)

Bond Formation

• Exothermic — energy released & decreases

  • Particles get closer, decreasing PE (energy needs to be released)

Relation to Overall Enthalpy Change of Reaction

Negative Enthalpy Change (exothermic)

• Energy released in bond formation > energy absorbed in bond breaking

Positive Enthalpy Change (endothermic)

• Energy released in bond formation < energy absorbed in bond breaking

Thermochemical Equations

A balanced chemical equation that includes the reaction’s enthalpy change (H). 

Example: CH4(g)+2O2(g)   →   CO2(g)+2H2O(l) H=-890.3kJmol

• Exothermic reaction because the overall enthalpy change is negative

H: amount of energy released (-) / absorbed (+) in a reaction when the number of 

moles of a compound indicated in the balanced chemical equation is used. 

• In the example: -890.3kJ1 mol CH4 or  -890.3kJ2 mol O2 

Designing & Analysing Calorimetry Experiments (Calorimetry Lab)

A chemical reaction’s enthalpy change (H) can be determined experimentally using a calorimeter (device used to measure heat changes in a system & its surroundings). 

We can design an experiment where the heat released/absorbed by a reaction causes a temperature change in the solution (usually water) within the calorimeter. 

Assuming no heat loss to the surroundings:

Qrxn=-Qsolution

where Qsolution=msolutioncsolutionTsolution

To find the enthalpy change of reaction (H in kJ/mol):

H=Qrxn (kJ)moles of the limiting reactant in the experiment (mol)

Calorimetry Lab

Metal w/ Higher Specific Heat

• Takes more heat energy to increase the temperature of metal

• Causes greatest increase in temperature of water in calorimeter

  • Releases more heat energy per degree → greater flow of heat energy into water → greater increase in water’s temperature 

• Causes more severe burn

  • Holds more heat at a specific temperature → more heat to transfer to skin → more severe burn

Mathematical Calculations

Heat Flow in Calorimetry Experiments

Heat is energy → measured in Joules (J)

Heat flow between the system and surroundings can be described as energy (J) being absorbed and released by a system. 

• Negative energy value: heat flowing out of the system

• Positive energy value: heat flowing into the system

Example problem

How much heat is lost when a 64.0g piece of copper cools from 375 ºC to 26ºC? The specific heat of copper is 0.38452 J/gºC. Place your answer in kJ. 

List of knowns: m=64.0g c=0.38452J/gºC T=26ºC-375ºC=-349ºC

Equation used: Q=mcT

Calculation: Q=(64.0g)(0.38452J/gºC)(-349ºC)1kJ1000J

=-8.59kJ

Answer: 8.59 kJ of heat is lost from the reaction (heat flows out of system)

Heat Absorbed/Released Given a Certain Amount of Substance

1. If given the amount of multiple reactants, find which reactant is limiting 

   1. Way 1: Compare the amounts of both reactants. The one with less quantity is the limiting reactant. 

   2. Way 2: Find the heat absorbed/released from the reaction using the given amount of each reactant. The reactant that releases/absorbs a smaller amount of heat is the limiting reactant. The already calculated amount of heat released/absorbed for that reactant is the answer. 

2. Calculate the heat absorbed/released in the reaction using the given amount of the limiting reactant 

Example Problem

How much heat will be released when 12.0g H2 reacts with 76.0g O2 to make water? (△H = -571.6 kJ)

Balanced Equation: 2H2 + O2  →  2H2O

Using Way 2: 

Find the heat absorbed/released from the reaction using the given amount of each reactant.

12.0gH21molH22.016g2-571.6 kJ2molH2=-1700 kJ

76.0gO21molH232.0gO2-571.6 kJ1molH2=-1360 kJ

76.0gO2 releases a smaller quantity of heat energy, so it is the limiting reactant 

Answer: 1360 kJ of heat will be released from the reaction. 

Objectives: 

• Define: entropy 

• Entropy changes during physical processes (melting, freezing, boiling, condensation, etc)

  • Describe the changes

  • Predict the sign of the entropy change Ssys

• Entropy changes for chemical reactions based on the number of moles & complexity 

  • Predict the sign of the entropy change Ssys

• Explain: enthalpy changes of a system influence the entropy change of the surroundings

• Second Law of Thermodynamics 

• Spontaneity of physical & chemical processes (related to entropy change for the universe)

  • Apply 2LT

  • Relate spontaneity to the natural tendency of systems to move towards disorder 

• Explain: endothermic processes can be spontaneous; exothermic processes can be non-spontaneous

Entropy

Symbol: S

Units: J/K (energy dispersion per temperature)

Conceptual Definition: A measure of randomness (chaos, disorder, etc)

• A dispersal of energy and matter increases randomness → more entropy 

Scientific Definition: Number of possible microstates 

Microstates—specific configurations of a system considering position & energy

• More possible microstates = more entropy 

Examples

1. Egg shell cracking

   1. Creates more objects → more microstates → more entropy

2. Small room with chair vs. Large room with chair

   1. Large room has more space (position) → more microstates 

Entropy Changes in Physical Processes 

Entropy of Various Phases

Solids (least)     →     Liquids     →     Gases (most)

• Gases have the most kinetic & potential energy → most entropy 

  • Arrangement of particles: furthest apart → high PE

  • Freedom of motion: greatest (almost no limitation) → high KE

Entropy of Physical Processes

• Melting & Vaporisation — increase in entropy (s→l & l→g)

  • Ssys > 0 (positive)

• Freezing & Condensation — decrease in entropy (l→s & g→l)

  • Ssys < 0 (negative)

Entropy Changes in Solutions

• Solid dissolving in a liquid — increase in entropy

  • More possible arrangements in aqueous solution → more entropy

• Gas dissolving into a liquid — decrease in entropy

  • More volume & possible arrangements in gas form → more entropy 

  • Less volume & space to move in liquid form → less entropy 

Entropy Changes in Chemical Reactions

Factors affecting Ssys where the chemical reaction is the system: 

1. Change in Moles of Gas (related to phase) — greatest influence 

   1. Increasing moles of gas increases entropy of system

   2. 2H2O(g) → 2H2(g)+O2(g) (2 moles of gas → 3 moles of gas)

      1. System’s entropy increases (Ssys>0)

2. Number of Moles — moderate influence

   1. Increasing number of moles increases entropy of system

3. Mixture vs. Non-mixture — moderate influence

   1. Mixtures (multiple types of molecules) have more entropy than non-mixtures

   2. C2H4(g)+H2(g) → C2H6(g) (mixture → non-mixture)

      1. System’s entropy decreases (Ssys<0)

4. Change in Phase Excluding Gas — moderate influence

   1. Liquids & aqueous solutions have more entropy than solids (more freedom of motion & greater space between molecules)

5. Complexity of Molecules — least influence

   1. More complex molecules have greater entropy because there is more vibration in the molecule and more possible microstates

   2. More complexity could refer to: larger molecules, more atoms involved, more electrons in the atoms, etc (e.g proteins are complex)

   3. HI(g) vs. HCl (g) (I has more electrons than Cl→more complexity)

      1. HI has greater entropy 

Enthalpy Change of System to Entropy Change of Surroundings

Hsys: enthalpy change of system Hsurr: enthalpy change of surroundings

Ssurr: entropy change of surroundings

Rule: 

Hsys is opposite in sign to Hsurr

Hsurr has the same sign as Ssurr

Explanation:

Since the space of the surroundings cannot be determined (always expanding & extremely large), energy is the main factor determining the surroundings’ entropy. Hence, the entropy change of the surroundings can be determined with the enthalpy change of the system. 

Hsys < 0     →     Hsurr > 0     →     Ssurr > 0 

• System lost energy. Surroundings gained energy. Surroundings increased in entropy. 

• E.g exothermic reactions

Hsys > 0     →     Hsurr < 0     →     Ssurr < 0

• System gained energy. Surroundings lost energy. Surroundings decreased in entropy. 

• E.g endothermic reactions

• Entropy of the system (Ssys) is not related and needs to be looked at separately.  

  • Can be determined using concepts of entropy change in physical processes and in chemical reactions mentioned above

Second Law of Thermodynamics

The entropy of the Universe is always increasing. The total change in entropy of the Universe is always positive. 

Spontaneity 

Definition:

A process is spontaneous if it agrees with 2LT by increasing the entropy of the Universe. 

A process is not spontaneous if it disagrees with 2LT by taking away from the entropy of the Universe. 

Spontaneous processes don’t require a constant source of energy. Once the process is started (activation energy is met), it will continue until completion. This is because systems have a natural tendency to move towards a state of greater disorder and energy dispersal. 

Example: combustion—once a fire is started, it will continue burning until completion

Equation: Suni=Ssys+Ssurr

Suni>0      —     process increases entropy of the Universe → spontaneous

Suni<0      —     process decreases entropy of the Universe → not spontaneous

• Suni refers to the observed process’ contribution to the entropy of the Universe (doesn’t refer to the overall change in entropy of the Universe, as it’s always positive)

• Ssys determined by applying concepts of entropy change in physical processes and chemical reactions to the examined system 

• Ssurr determined with the system’s enthalpy change (e.g exothermic/endothermic)

Endothermic & Exothermic Processes

Endothermic: system gains energy

• Surroundings lose energy. Entropy of surroundings decreases (Ssurr<0).

• Indicated by surroundings getting colder

• Suni is not necessarily negative (not spontaneous) because Ssys could be a greater positive than the negative Ssurr, making Suni positive 

Exothermic: system loses energy

• Surroundings gain energy. Entropy of surroundings increases (Ssurr>0). 

• Indicated by surroundings getting hotter 

• Suni is not necessarily positive (spontaneous) because Ssys could be a greater negative than the positive Ssurr, making Suni negative 

System-Surrounding Disagreement 

Suni=Ssys+Ssurr

If Ssys and Ssurr are opposite in sign, the sign of Suni, which determines spontaneity, depends on the temperature of the system. You can only change the temperature of the system, not the surroundings. 

High Temp: magnitude of Ssys>Ssurr

• If Ssys < 0 and Ssurr > 0, the negative Ssys would outweigh the positive, making Suni negative

• Vice versa if Ssys > 0

Low Temp: magnitude of Ssys<Ssurr

• If Ssys < 0 and Ssurr > 0, the positive Ssurr would outweigh the negative Ssys, making Suni positive

• Vice versa if Ssys > 0 

Spontaneity Table