Chem Quiz 1 Review

Version 1



 

Version 2

 

Physical and Chemical Properties and Changes

What is Matter?

 

Properties

 

Types of Properties

  1. Chemical Properties
  2. Physical Properties

 

Chemical & Physical Properties

–         Includes qualitative and quantitative properties

 

Qualitative vs. Quantitative Properties

–        Example- The car is red.

–        Example – It is 5°C outside.

 

Chemical vs. Physical Change

Examples

Chemical Physical

Burning wood                         Sawing wood

Decomposing water                Freezing water

Nail rusting                             Cleaning rust with steel wool

 

Chemical Change Clues

  1. A gas is given off.
  2. A color change occurs.
  3. A precipitate (solid) is formed.
  4. Heat is absorbed or given off.
  5. Electrons are transferred.

If the change is difficult to reverse, it hints at a chemical change.

Keep in mind these are only clues that a chemical change may have taken place.  Proof still depends on a new substance being formed with new properties.

 

 

Periodic Table and Bohr-Rutherford Diagram

 

Structure of an Atom

 

What is the structure of an atom?

Nucleus – center of the atom

¡  Consists of Neutrons and Protons

¡  Proton

n  Has a positive (+) charge

n  Has a relative mass of 1

n  Protons not removed or added to atoms, to give them charge

n  Determines the atomic number

n  Found inside the nucleus

n  Because they are heavy and inside the nucleus, protons only leave the nucleus in “nuclear reactions”

 

¡  Neutron

n  Has no (0) charge

n  Has a relative mass of 1(roughly same mass as a proton)

n  Found inside the nucleus

n  Because they are heavy and in the nucleus, neutrons only leave atoms in nuclear reactions

 

Electron

¡  Has a negative (-) charge

¡  Has a relative mass of 0

¡  Determines the ion

¡  Found outside the nucleus

¡  Smallest subatomic particle

¡  About 1 /1837th the mass of a proton or neutron

¡  Since they are light and not inside the nucleus, electrons can be removed or added to atom, causing atom to become charged

 

 

 

 

 

 

 

 

 

 

 

 

 

Periodic Table of Elements

 

 

 

 

 

 

 

 

The Current Periodic Table

n  The elements are arranged by increasing ATOMIC NUMBER

n  The horizontal rows are called periods and are labeled from 1 to 7.

n  The vertical columns are called groups are labeled from 1 to 18.

 

Inside an Element Square

n  Different periodic tables can include different information, usually:

¡  atomic number

¡  element symbol

¡  atomic mass

¡  State of matter at room temperature.

Atomic Number

n  This refers to how many protons an atom of that element has

n  No two elements have the same number of protons

n  If they did they would be the same element

Atomic Mass

n  Atomic Mass refers to the “weight” of the atom.

n  It is derived at by adding the number of protons with the number of neutrons.

n  Atomic mass = (P+) + N0

 

Isotopes

¡  Isotopes are two of the same element with different masses/different number of neutrons

¡  Have same number of protons, but different number of neutrons

¡  It is atomic number (number of protons) that identifies an element

¡  Atomic mass on periodic tables accounts for all forms of isotopes of the element

Valence Electrons

n  The number of valence electrons an atom has may also appear in a square.

n  Valence electrons are the electrons in the outer energy level (orbit) of an atom.

n  Valence electrons are transferred or shared when atoms bond

Properties of Periods and Groups

 

n  Sizes of the atoms generally decrease as we move from left to right across a period

n  As you move from left to right across a period, the ability of the atom to attract another electron increases

n  This property is called electronegativity

Hydrogen

n  The hydrogen square sits atop Group 1, but it is not a member of that family. Hydrogen is in a class of its own.

n  It’s a gas at room temperature.

n  It has one proton and one electron in its one and only orbit.

Alkali Metals

n  The alkali family is found in the first column of the periodic table.

n  Atoms of the alkali metals have a single electron in their outermost orbit, in other words, 1 valence electron.

n  They are shiny, have the consistency of clay, and are easily cut with a knife.

n  They are the most reactive metals.

n  React violently with water.

n  Alkali metals are never found as free elements in nature.

Alkaline Earth Metals

n  They are never found uncombined in nature.

n  They have 2 valence electrons.

n  Alkaline earth metals include magnesium and calcium, among others.

Transition Metals

n  Transition Elements

n  These are the metals you are probably most familiar: copper, tin, zinc, iron, nickel, gold, and silver.

n  Good conductors of heat and electricity.

Halogen Family

n  The elements in this family are fluorine, chlorine, bromine, iodine, and astatine.

n  Halogens have 7 valence electrons, which explains why they are the most active non-metals. They are never found free in nature.

Noble Gases

n  Noble Gases are colorless gases that are extremely unreactive because their outermost orbit is full.

n  Because they do not readily combine with other elements to form compounds, the noble gases are called inert.

n  The family of noble gases includes helium, neon, argon, krypton, xenon, and radon.

n  All the noble gases are found in small amounts in the earth's atmosphere.

Rare Earth Elements

n  The thirty rare earth elements are composed of the lanthanide and actinide series.

 

 

The Bohr Model of the Atom

n  An effective way to represent the first 20 elements.

n  Each electron orbit is shown as a ring around the nucleus

 

n  The 1st orbit will hold a maximum of 2 electrons

n  The 2nd orbit will hold a maximum of 8 electrons

n  The 3rd orbit will hold a maximum of 8 electrons

Example:

 

 

Bohr-Rutherford Diagrams

1)      Using periodic table, find the number of protons and neutrons

Protons = Atomic Number

Neutrons = Atomic Mass – Atomic Number

2)      Put values of protons and neutrons in nucleus of atom

Example – Oxygen

P = Atomic Number

= 8

N = Atomic Mass – Atomic Number

= 16 – 8

= 8

 

1) Determine the number of electrons

(= number of protons = atomic number)

Determine number of electrons from charge

If atom is neutral (electrons = protons = atomic number)

Add electrons to the shell, and once one shell fills, start another

Remember the maximum number of electrons for the shells is 2, 8, 8, 8…..

 

2) Place electrons in their proper orbit (2, 8, 8)

 

 

What does it mean to be reactive?

n  We will be describing elements according to their reactivity.

n  Elements that are reactive bond easily with other elements to make compounds.

What makes an element reactive?

¡  An incomplete valence electron level (outer orbit).

¡  All atoms (except hydrogen and helium) want to have 8 electrons in their very outermost energy level (this is called the rule of octet.)

¡  Atoms bond until this level is complete. Atoms with few valence electrons lose them during bonding. Atoms with 5, 6, 7, or 8 valence electrons gain electrons during bonding.

 

 

 

 

 

Lewis Structures

n  Valence electrons involved when atoms are bonding/ forming compounds

n  Valence electrons are electrons in the outermost orbit in an atom

n  Lewis structures only show valence electrons of an element – much easier

Rules

n  Element symbol placed in the center (represents Nucleus – protons, neutrons, and inner electrons (not valence electrons))

n  Symbol assumed to have 4 sides, and valence electrons placed as dots around the four sides

n  Dots placed singly, then paired, in order from (N S E W)

 

Ions

 

 

To find out what charge an ion will have, must look at Bohr-Rutherford diagram of an element, and determine whether that element will gain or lose electrons to have full outer orbit/energy level

Atom does whatever is easier – if it has between 1 – 3 electrons, loses them; if 5-7 electrons, gains remaining electrons

Ions are only in group 1, 2, 13, 15, 16, 17

Carbon group does not have ions or ionic bonding

To draw ions, draw Bohr-Rutherford diagram, put diagram in square brackets, place charge on top right of bracket

To name positive ion – name is the same as the name of the element followed by “ion”

To name negative ion – name is determined by removing the end, adding “ide”, than writing “ion”

Ex. Oxygen ion = oxide ion, Phosphorus ion = Phosphide ion

 

How does an atom become charged?

 

Multivalent Ions

Ex. Cu+4 = Copper (IV) ion, Ni+3 = Nickel (III) ion

 

 

Ionic Compounds

Recall Ions:

-          Elements on the left of the Periodic Table (metals) have “+” charge because they lose electrons (become more positive/less negative)

-          Elements on the right of the Periodic Table (non-metals) have  “ – negative” charge because they gain electrons (become less positive/more negative)

Ex.

-          (x)+1 + (y)-1 = (+1)+(-1) = 0, the charge is neutral or 0

 

Ex.

Aluminum Oxide Continued

Having lost 3 electrons, the aluminum has a charge of +3.

Having gained 2 electrons, the oxygen has a charge of -2.

 

Naming Ionic Compounds

  1. 1. Write the name of the metal
  2. 2. Write the name of the non-metal, changing the end of the non-metal name to “ide”

Ex. Sodium chloride, hydrogen fluoride

 

Writing Formulas for Ionic Compounds

The trick to finding the ratio of atoms in the molecule

****** Those

  1. Number the groups 1-3 going from left to right (skip the Transition Metals)
  2. Number the groups 1-3 going from right to left  (skip the Nobel Gases)

 

Then …

Ex.

This becomes: (Al2O3)

 

 

 

Multivalent Compounds

Naming

e.g. Write the name for the following compounds,

CuF

Copper (I) Fluoride

PbI2

Lead (II) Iodide

CaF2

Calcium Fluoride

-Calcium is not a transition metal, so it does not use a roman numeral

 

-          Do the reverse of the criss cross method

-          Take the subscript and use it as the charge on the opposite symbol

Ex.

 

IMPORTANT

Ex.

CuS = Copper (I) Sulphide

–        Indicates that S has a charge of -1

–        From Periodic Table, we know charge of Sulphur is -2

–        So, must multiply both charges by 2

–        This leads to:

CuS = Copper (II) Sulphide

 

 

Ionic Compounds

 

 

Properties of Ionic Compounds