Grade 11 Chem Unit Test #1

Physical Properties—Qualitative

-State

-Colour

-Clarity—transparent, translucent, opaque

-Viscosity

-Odour

-Taste

-Crystal, powder, granular

-Shiny/dull (lustrous)

-Ductile

-Malleable

-Texture

-Mixture

Physical Properties—Quanitative

-Volume

-Mass

-Density

-Temperature

-pH

-Viscosity

-Conductivity

Chemical Properties

-react with acid

-react with base

-combustibility (react with O2)

-rust/ oxide (react with O2)

Physical Changes

-change of state (phase change)

-dissolving

-form change

E.x., evaporation, salt + water, bending lead

Chemical Changes

-heat/ light released/ absorbed

-colour change

-precipitate (new solid)

-gas (bubbles, vapors, new smell)

-can’t reverse

Groups

Vertical columns

Periods

Horizontal rows

What are diatomic molecules? Identify seven of these.

Ions with more than one atom

H2 O2 Br2 F2 I2 N2 Cl2

Acids

pH value: 1-6 or <7

Colour in phenolphthalein: colourless

Colour of litmus paper: red

Ions present in solution: H+

Physical properties: grippy, sour, corrosive, conductive, aqueous

Bases

pH value: 8-14 or >7

Colour in phenolphthalein: pink

Colour of litmus paper: blue

Ions present in solution: OH-

Physical properties: bitter, corrosive, conductive, aqueous, slippery

Atom

The smallest particle of an element

Atomic Model & Theory Timeline

Democritus: he proposed that matter could not be divided into smaller pieces forever. He claimed that matter was made of small, hard particles that he called “atoms”.

Dalton: he created the very first atomic theory. Dalton viewed atoms as tiny, solid balls.

Thomson: he showed that the atom was made of even smaller things. His atomic model was known as the “raisin bun model”

Rutherford: he discovered protons and the nucleus. He showed that atoms have (+) particles in the center, and are mostly empty space.

Bohr: he improved on Rutherford’s model. He proposed that electrons move around the nucleus in specific layers, or shells.

Chadwick: he discovered neutrons. Working with Rutherford, he discovered particles with no charge; these particles were called as neutrons.

Modern: work done since 1920 has changed the model. The new atomic model has electrons moving around the neon a cloud.

Democritus

The smallest bit of matter was an atom

Dalton

Described all matter as being composed of tiny particles called atoms, represented by different sized spheres

Thomson

“Plum pudding” in which a positively charged sphere had negatively charged particles embedded in it

Rutherford

Planetary model in which all the positive charge and most of the mass of the atom is located in the center of atom, with the electrons orbiting the nucleus

Bohr

Modified planetary model, stating electrons have specific amounts of energy and occupy orbits, or energy level

Quantum mechanical model

Determined the electrons not found in precise orbit, but existed in electron clouds called orbitals

Electron

Location in the atom: in orbit

Mass: 1/2000 u

Charge: -

Proton

Location in atom: in nucleus

Mass: 1 u

Charge: +

Neutron

Location in atom: in nucleus

Mass: 1 u

Charge: o

Chemical Notation

A—mass # (p+, n0)

X—chemical symbol

Z—atomic # (p+, e-)

Isotope

A form of a chemical element where atoms have the same number of protons but a different number of neutrons

Ion

When an atom loses or gains electrons and becomes charged

Atomic mass (average)

Average of all isotopes

Average atomic mass formula

E.g., Mav = (Mir-191)(%ir-191)+(Mir-193)(%ir-193)

(%ir-191) = 100% - (%ir-193)

Cations

Positive ions

Metals tend to form cations

Anion

Negative ions

Non-metals tend to form anions

Brought to you by…

Dimitri Mendeleev —

A Russian Chemist who discovered many elements and developed the modern Periodic Table of Elements

Alkali Metals

Group 1 on periodic table

Very reactive

Can explode when they come in contact with water or oxygen.

They are not found in nature on their own — they are a part of a compound

Easily lose an electron to become a positive ion.

Alkaline Earth Metals

Group 2 on the periodic table.

Semi-reactive

Also, easily lose an electron to become positive ions

Found readily in nature

Transition Metals

Group 3-12 on the periodic table

Conduct heat and electricity

Ductile — able to be deformed without losing toughness!

Malleable, bendable, flexible, etc.

Halogens

Group 17 on the periodic table

The word Halogen means salt forming

Highly reactive with the Alkaline Earth and Alkali Metals

Noble Gases

Group 18 on the periodic table

Glow when you pass an electric current through the gas!!

Doesn’t react with other elements!

That’s why it’s relatively safe to inhale helium - except you can suffocate because you’re not providing your brain with any oxygen…

Non Metals

Hydrogen, Carbon, Nitrogen, oxygen, phosphorus l, sulfur, + selenium!

These are dull, brittle, and reactive with other elements.

Technically, everything above this bolded stair-like line is a nonmetal. However, only the above seven are in the official nonmetal family

Poor Metals

Between the transitional metals and the metalloids

Just as their name suggests these have poor qualities: soft, not great conductors, and lightweight

Metalloids

Boron, silicon, germanium, arsenic, antimony, tellurium + polonium!

Conducts electricity and essential for the production of electronics and computer chips!!

AKA: the stair stepper

Lanthanides

Elements with an atomic number of 57-71

Soft, malleable, and highly conductive

Usually found as alloys (compounds with multiple metals)

Actinides

Elements with an atomic number 89-103

Only the first four (Ac, Th, Pa, & U) occur naturally

All radioactive!! — spontaneous emission of radiation (a type of wave) due to an unstable nucleus

Making a new element

Every element after uranium is synthetic (man made)

High energy particle accelerator: machines that move atomic nuclei at high speeds until they collide

How is the periodic table organized

Firstly, by number of protons

Mass too!

As you go from left to right, each element is larger than the one before it

States of matter and temperature

All elements are shown at room temperature (25 C)

Most elements (103) are solids, 11 are gases, and 2 are liquids

Metals throughout the table

Most metals are found on the left hand side and middle of the periodic table

Gasses are found on the right hand side

Valence electrons

A Valence electron is an electron in the outermost shell or energy level of the atom

The 8 major columns depict the same number of valence electrons as their column number

Note: the 8th column has full shells!

Atomic Radius

The distance from the center of an atom to the boundary within which the electrons spend 90% of their time

Can be determined using x-ray crystallography, neutron diffraction, or electron diffraction

The radius of an atom is the distance from nucleus to the outermost orbit

Atomic size increases down a group due to the increasing number of orbits occupied and a decrease in effective nuclear charge (the apparent charge experienced by valence electrons)

Atomic size decreases across a row due to increasing number of protons in the nucleus that exert a stronger effective nuclear charge on valence electrons

Ionization Energy

The amount of energy required to remove the outermost electron from an atom or ion in the gaseous state

When an atom loses an electron the remaining ion is positively charged. Energy is required to remove an electron

A(g) + energy — A+(g) + e-

After one electron is removed, it is still possible to remove more.

A+(g) + energy — A2+ + e-

More energy is required for second ionization because there are now less electrons but the same number of positive charges attracting them to the nucleus

Noble gases require the most ionization energy. They are small stable elements. Tight hold on electrons, high Zerg.

Ionization energy decreases down a group as the outer electrons are held further away from the nucleus and require less energy to be removed

Ionization energy increases across a row as the effective nuclear charge becomes larger as more protons are added but the number of orbits stays the same

Electron Affinity

The energy absorbed or released when an electron is added to a neutral atom

If a neutral atom gains electrons it may become negatively charged. In this process energy may be released, or energy may be needed to add the electron

This energy is the atom’s electron affinity — how likely it is to gain an e-

Negative values indicate energy is released

The more negative a value, the more stable the ion - more likely to gain e-

Electron affinity decreases down a group

Electron affinity increases across a period

Zeff

These trends can all be linked to electron arrangement and Effective Nuclear Charge “pull” (Zeff), which is the net positive charge experienced by an electron in a multi-electron atom

Zeff increases from left to right across a period and decreases down a group on the periodic table

Protons vs. Electrons

In a multi-electron atom, the positive charge of the nucleus is partially offset by the negative charge of the inner-shell electrons

Shielding Effect

These inner-shell electrons (core electrons) “shield” the outer electrons from the full attractive force of the positively charged nucleus

Chemical Bonding

Binding is caused by the interactions between the valence e- of atoms

Atoms tend to gain stability (ie. a lower energy state) by gaining, losing or sharing valence electrons, in order to gain a complete octet

Properties of Ionic vs. Covalent Compounds

If a metal atom is bonding with a non-metal atom in a compound, then the bond is ionic

Example: NaCl, K2SO4, MgCl2

If all the elements in the substance are non-metals, then the bond is covalent

Example: C2H5OH, C12H22O11, H2O, CO2

If all the elements in the substance are metals, then the bond is metallic

Example: alloy - bronze, steel, brass

Ionic

Soluble in H2O

Electrolytes (conduct in H2O)

Solids at room temperature

Higher fixed points (melting or boiling point)

Covalent

Varying solubility in H2O

Not electrolytes

Solid, liquid or gas at room temp

Low fixed points (melting or boiling points)

Electronegativity (EN)

Is the ability of an atom to attract electrons when bonded to another atom

Electronegativities are listed under the atomic number of your periodic table.

Across a period: EN increases due to a smaller radius and higher effective nuclear charge (“pull”)

Down a group: EN decreases due to larger size and weaker effective nuclear charge

The electronegativity difference (triangle EN) between two atoms can be calculated in order to predict the type of bond that would form between them

If the change in EN >1.7, the bond is IONIC. If the change in EN <1.7, the bond is COVALENT

The Bonding Continuum

Bonding between atoms can be classified on a continuum, which depends on the change in EN between the two atoms

Pure covalent bonding

H2

Change in EN = 0-0.4

Covalent bond, two shared electrons

air. The electrons of both atoms spend an equal amount of time circling BOTH hydrogen nuclei

Seven elements exist as diatomic molecules- two of the same atoms joined by a pure covalent bond

Polar Covalent Bonding

HCl

Change in EN 0.4<EN<1.7

Cl is more electronegative, so it attracts electrons more than H. Thus, the chlorine “side” of the molecule is more negative and the hydrogen “side” of the molecule is more positive.

Dipole

We say that the molecule has a dipole and show it with an arrow pointing to its negative side.

What kind of substances would water be attracted to (adhere to)?

Charged object

Explain how capillary action would work in a tree, where water is drawn up from from the roots to the leaves against the force of gravity.

Cohesion - H2O sticks to itself

Polar covalent bond

A type of chemical bond where electrons are shared unevenly between two atoms, creating a partial negative charge on the more electronegative atom and a partial positive charge on the less electronegative atom

(EN diff. <1.67 and >0.40)

Nonpolar covalent bond

A chemical bond formed when two atoms share electrons equally because they have very similar or identical electronegativity

(EN difference < or equal 0.40)

Dipole-dipole forces

Each molecule is known as a dipole. The attraction between the positive end of one dipole and the negative end of another is called dipole-dipole force

Dipole-dipole forces arise between polar molecules

Dipole induced dipole force

When a polar molecule approaches a Nonpolar molecule, the electron cloud of the Nonpolar molecule may become distorted, causing the Nonpolar molecule to become temporarily polar. For example, if the negative end of the polar molecule approaches the Nonpolar molecule, the electron cloud of the Nonpolar molecule will be repelled, causing a slight positive charge at that end of the Nonpolar molecule. The resulting attractive force is called a dipole-induced dipole force

Dipole-induced dipole forces arise between polar and Nonpolar molecules

London dispersion forces

Even when the molecules are Nonpolar, random variations in the distribution of electrons can cause parts of these molecules to become slightly charged. This imbalance leads to very tiny, short-lived attractions between molecules called London dispersion forces

London dispersion forces arise between Nonpolar molecules

Polar molecule

Has an uneven distribution of electron density, resulting in one end of the molecule being slightly positive and the other end being slightly negative

Nonpolar molecule

Has an even distribution of electrons density, resulting in no separation of positive and negative charges and a net dipole moment of zero

(Symmetrical)