Ch2 Components of Matter

Element – simplest type of matter with unique physical and chemical properties (only one kind of atom)

Molecule – independent structural unit consisting of two or more atoms that are chemically bound together

Compound – composed of two or more different elements that are chemically bound together

Mixture – a group of two or more substances (elements and/or compounds) that are physically intermingled.

Laws

Mass Conservation – the total mass of substances does not change during a chemical reaction.

Definite Composition – a particular chemical compound is composed of the same elements in the same

parts (fractions) by mass.

Multiple Proportions – if elements A and B react and form two different compounds, the different masses of B

that combine with a fixed mass A can be expressed by a ratio of small whole numbers.

Dalton’s Atomic Theory and its explanation of Mass Laws

Read postulates 1-4, page 46

Mass Conservation – Atoms cannot be created or destroyed, cannot be converted into other types of atoms, a

chemical reaction where the only difference is the combination of atoms cannot produce a

mass change.

Definite Composition – A compound has a specific ratio of atoms, so each elements has a fixed fraction of mass

Multiple Proportions – Atoms of same element have same mass, so if different masses of element B combine

with a fixed amount of A, then the resulting compound will differ by a small whole

number ratio.

Nuclear Atom Model

Cathode Ray Tube experiment – J.J. Thoompson

Millikan Oil Drop experiment

Rutherford’s a - scattering experiment

Atomic Theory Today

Atom – “electrically neutral, spherical entity composed of a positively charged central nucleus surrounded by one or more negatively charged electrons.”

Nucleus – consists of protons and neutrons (except for hydrogen).

Atomic Number (Z) – number of protons, “each element has a different atomic number from that of any element”

Mass Number (A) – total number of neutrons and protons in atom (number of neutrons = A – Z)

Atomic Symbol (X) – symbol of element (examples C- carbon, O – oxygen, Na – sodium, etc…)

Isotopes – A particular element have different mass numbers due to the number of neutrons it contains in the nucleus. Sample Problem 2.2 (pg 53)

Read “A Modern Assessment of the Atomic Theory” pgs 55 – 56

Periodic Table

Figure 2.11

Increasing Atomic Number as you go from left to right and top to bottom.

Period – rows of periodic table

Groups – columns of period table

Main group elements – The eight A groups

Transition Elements – The ten B groups

Types of elements – Metals, Non-metals, Metalloids

Bonding

Ionic compounds – transfer of electrons from one element to another. Typical when non-metal and metal react.

Cation – positively charged ion

Anion – negatively charged ion

Neutral – no net charge

Noble gas relationship – metals tend to lose electrons and non-metals tend to gain electrons to “form

ions with the same number of electrons as in the nearest noble gas.” Fig. 2.14, 2.15

Covalent compounds – sharing electrons between elements, usually occurs when a non-metal reacts with a non-metal. “Most covalent substances consist of molecules, whereas no molecules exist in a sample of an ionic compound.”

Chemical Formulas

Empirical Formula – relative number of elements in compound

Molecular Formula – actual number of elements in compound

Structural Formula- shows number of atoms and the types of bonds between them.

Naming

read pgs 64 – 71

Molecular Mass

Sum of Atomic masses of all elements in molecule (or compound). Sample Problem 2.13 (pg 72)

Ch3

Mole – SI unit for amount of substance. Number of entities related to mole by Avogadro’s number

(N_A = 6.022 x 10²³ entities)

Molar Mass – mass per mole (g/mole)

Look at pg 90 and sample problem 3.1 to convert between mass, moles, and entities

Mass percent of element =

Look at Sample Problem 3.4,3.5, and 3.6 to obtain empirical formulas and molecular formulas of unknown compounds.

Balancing Chemical Equations – the same number of each type of atom must appear on both sides of the chemical equation (products and reactants).

Calculating Amounts of Reactants and Products pg 106 gives a good explanation

Limiting Reagent – “the reactant that yields the lower amount of product.”

Percent yields =

Problem 69 –

Solid Iodine trichloride (ICl₃) is prepared by reaction between solid Iodine (I_2(s) ) and gaseous chlorine (Cl_2(g)) to form monochloride crystals (ICl_(s)), followed by treatment with additional chlorine (Cl_2(g)).

a) Write a balance equation for each step.

b) Write an overall balance equation for the formation of iodine trichloride.

c) How many grams of Iodine are needed to prepare 31.4 kg of final product?

Step 1) I_2(s) + Cl_2(g) à 2ICl_(s)

Step 2) ICl_(s) + Cl_2(g) à ICl_3(s)

b) We want to show only initial Reactants going to final Products:

I_2(s) + Cl_2(g) à ICl_3(s)

Need to somehow balance and cancel out the intermediate products. Add step 1) and step 2). Remember to subtract the compounds that appear on different sides of the equations. By multiplying Step 2) by 2 (to cancel out ICl₃) we get:

I_2(s) + Cl_2(g) à 2ICl_(s)

2[ICl_(s) + Cl_2(g) à ICl_3(s)]

----------------------------------

I_2(s) + 3Cl_2(g) à 2ICl_3(s)

Check that the chemical equation is balanced:

Reactants Products

I 2 2

Cl 6 6

Molecular Mass of ICl₃ :

I – 126.9

Cl – 3x 35.45

-----------------

233.25 g/mol ICl₃

Stoichiometry

Solute – smallest amount in the solution

Solvent – largest amount in the solution

Molarity (concentration) M– in units of moles of solute per Liter of solution

Dilution

M₁V₁ = M₂V₂ , where M₁ is the higher concentration (starting value) and M₂ is the diluted solution, V₁ and V₂ are their respective volumes.

Ch 6 Thermochemistry

Energy

Law of Energy Conservation – “the total energy of the universe is constant.”

DE_universe = DE_system + DE_surrounding = 0

System – part of universe we are going to observe

Surroundings – everything else relevant to the change (except for system).

Change in energy: final – initial = Product – Reactants

Heat – thermal energy, q.

Work – “all other forms of energy transfer” (mechanical, electrical, etc…), w.

Heat

Flowing out of system, -q => -DE

Flowing into a system , +q => +DE

Units

Joule (J) = kg m² s^-2, 4.184 J = 1 cal

State Functions

Property only dependent on the current state of the system, not on the path system took to reach state.

Enthalpy

Internal Energy plus product of Pressure (P) and Volume (V)

Change in Enthalpy equals the heat gained or lost (by the system) at constant pressure

Endothermic – absorbs heat, increase in enthalpy of the system

Exothermic – release of heat, decrease of enthalpy of the system

Types

Heat of Combustion – Enthalpy change in Combustion reaction

Heat of formation – Enthalpy change to form a compound from its elements

Heat of fusion – Enthalpy change of the physical process of melting

Heat of vaporization – Enthalpy change of the physical process of vaporization (boiling)

Calorimetry

Specific heat Capacity (c) – Energy required to raise Temperature by 1^oC of 1g of pure substance (J/g^oC)

Heat capacity – Energy required to raise Temperature by 1^oC of entire substance (J/^oC)

q= m c DT (m –mass, c – specific heat capacity, DT – change in Temperature)

Sample Problems 6.4, 6.5 (pgs 234-236)

Hess’s Law of Heat of Summation

The overall process has a final Enthalpy change that can be summed up by the individual Enthalpy changes for each step.

(sample problem 6.9, chapter problems6.81,86,95)

Ch 7

Matter and Energy

Matter and Energy are related, and both have properties of waves and particles (wave-particle duality).

Property Examples: Light (Energy, classically thought of as a wave) – wave nature (diffraction pattern), also particle nature (light is made up of photons)

Electron (classically thought of as matter) – wave nature (has wave-like orbitals), also particle nature (localized electron density)

Equations relating matter to energy:

E – energy, m – mass, c – speed of light, n – energy level, h – planck’s constant, - frequency, l - wavelength

Electromagnetic Spectrum

Wave Nature of Light

Three main Characteristics:

1. Wavelength (l) – length of 1 cycle

2. frequency (n) - # cycles/sec

3. Speed (meters/sec) of traveling wave

Picture

2 complete cycles are shown.

The frequency is –

ß------------------2 seconds ---------------------à

You can relate the wavelength and frequency to the speed of the traveling wave:

Analysis of units -

Electromagnetic Radiation (light) travels at c – 3.00 x 10⁸ m/s (speed of light), so for light:

exercise

A dental hygienist uses x-rays (l = 1.00 x 10^—10m) to take a series of dental radiographs while the patient listens to a radio station (l =325 cm) and looks out the window at the blue sky (l =473 nm).

What is the frequency (in s^—1) of each electromagnetic source?

rearrange to obtain frequency

for the x-rays

for the radio

for the blue light

Quantum Theory

1. Energy is quantized ( ), so energy of a photon has this equation,. Remember that the frequency is related to the wavelength by the speed of light, so we can also find the energy if we do not know the frequency, only the wavelength:

substitute in the equation for energy of a photon,

For an atom in order to “promote” an electron from a low state n₁ (example n=1) to a higher state n₂, we need to figure how much energy we need the atom to absorb in order to make this happen:

Once we know that amount of energy, how are we going to put the energy into the atom for this “promotion” to occur? Remember photons are related to electrons, so we can say:

This means we need a photon with a high enough energy (or wavelength) to make the “promotion” happen:

rearranging, we can figure out what wavelength is needed

2. Since energy is quantized, and matter is related to energy, we can assign a wavelength to any object with the deBroglie wavelength equation:

where h – planck’s constant, m – mass, - speed

exercise

Find the deBroglie wavelength of an electron with a speed of 1.00 x 10⁶ m/s. (electron mass – 9.11 x 10^—31kg, h – 6.626 x 10^—34 kg m²/s)

Other useful equations:

(R- Rydberg constant) “This is for the atomic spectra of hydrogen.”

For atoms like Hydrogen (1 electron atoms, Li2+, Be3+, He+)

,

(Remember we are only dealing here with only one type of atom (or a mole of the same atoms) in these equations, so our Z (atomic number) value will be the same for both n_final and n_initial.)

Ionization Energy – energy required to take an electron from a low state to (remove the electron from the atom). The Energy at is designated as 0. For hydrogen:

since

where

(two negatives give a positive)

And removing the negative value on the energy we have:

Look at the Energy States of the Hydrogen Atom (pg 264 Silberberg)

Terms

Ground state – lowest lying state (for atoms like hydrogen n=1)

Excited state – n > 1

Heisenberg Uncertainty Principle

States that we cannot know exactly the position and the velocity of a particle at the same time.

- uncertainty in position, - uncertainty in speed, m – mass

Quantum Numbers

Look at this section in the book (pgs 275 – 277) – Remember this is not a complete set, you will have another quantum number in Ch8, m_s.

Name	Symbol	Values
Principal quantum number	n	Positive values (1,2,3,4,….)
Sublevel	l	0….n-1
Orbitals	ml	-l….0…+l
Spin	ms	-1/2, +1/2

Exercises

An alpha particle (m – 6.6x10^—24g) emitted by radium travels at 3.4x10⁷ 0.1x10⁷ mi/h.

Find the deBroglie wavelength, and the uncertainty in its position.

In firework shows, Light of a characteristic wavelength is related to the presence of a particular element. What are the frequency and color of light for each of the following elements?

Li⁺(671 nm), Cs⁺ (456 nm), Ca²⁺ (649 nm), Na⁺ (589 nm)

For colors look at Electromagnetic spectrum (Visible light portion). For frequency

Li⁺(red) Cs⁺(blue)

Ca²⁺(orange-red) Na⁺ (yellow-orange)

Ch8

Many Electron Atoms

We can designate 4 quantum numbers to describe electrons in an atom. Table 8.2 Silberberg, or above.

On our way to describing the periodic table there are certain rules we must follow:

Pauli Exclusion Principle – no two electrons in the same atom can have the same four quantum numbers, so each electron has a unique set of quantum numbers.

Exercise

Use the periodic table to identify the element with the electron configuration 1s²2s²2p⁴. Write the orbital diagram, and give the quantum numbers of its sixth electron.

Z= 8 Oxygen. For sixth e^—n = 2, l = 1 m_l = 0 m_s = +1/2

Electrostatic Effects & Splitting of Energy Levels – electrons in the ground state occupy the orbitals of lowest energy. In Ch 7, we only had one electron in the atom, now we have many, so there will be repulsion forces between the electrons. This repulsion causes the splitting of energy levels. In CH 7 we only dealt with the n state energy levels, know we deal with n state and the sublevel l .(i.e – n=3, can be broken down into 3s,3p,3d)

Factors:

(1) The farther the opposing charge, the weaker the interaction.

(2) The higher the opposite charge values, the stronger the interaction.

(3) Like charges repel

Look at figures on pg. 293. All of these factors will determine the effective nuclear charge (Z_eff) each electron “feels” in a particular orbital.

Build-up Principle (Aufbau) – To determine electron configuration of the elements, start at the beginning of the periodic table and add one electron per element (as well as proton/neutrons) to the lowest energy orbital available. This will result in the ground state electron configuration of the atom. (Look at pgs. 295 – 302 for rules and exceptions).

Hund’s Rule – When filling orbitals, remember that “when orbitals of equal energy are available, the electron configuration of lowest energy HAS the maximum number of UNPAIRED electrons with parallel spins.

Major Connections – Orbitals are filled in order of increasing energy (slight exceptions), which lead to outer electron configurations that recur periodically, which lead to chemical properties that occur periodically.

Categories of electrons

Look at definitions on pg 302 (Inner, Outer, Valence).

Trends of Periodic Table

Pgs 304 -311.

Atomic size (transition metals’ sizes stay somewhat constant)

Ionization Energy (opposite of atomic size)

Electron Affinity (opposite of atomic size)

Metallic Behavior

Ion size vs Atomic Size

Acid – Base behavior in oxides

Isoelectronic configurations

Magnetic Properties (paramagnetic – unpaired electrons, diamagnetic – paired electrons)