Chemistry Ch1-3

Chapter 1: Atomic Structure and the Periodic Table (P.T.)

In an atom, the bigger number is the mass number and the smaller is the atomic number.

• Atomic Number: The number of protons in an atom (in the nucleus).

• Mass Number: The sum of the protons and neutrons in an atom.

Remarks:

1. The number of protons equals the number of electrons in a neutral atom.

2. Number of neutrons = Mass Number – Atomic Number.

3. Electrons orbit the nucleus (outside the nucleus) in shells (energy levels). This is similar to the solar system analogy.

Exercises:

• Draw the Bohr-Rutherford (B-R) of:

  • 1632S: p: 16, n: 16, e: 16 (2, 8, 6)

  • 3919K: p: 19, n: 20, e: 19 (2, 8, 8, 1)

Atoms can lose or gain electrons to become an ion. An ion is a charged particle (not neutral).

• If an atom loses electron(s), it becomes a positive ion (cation).

• If an atom gains electron(s), it becomes a negative ion (anion).

Isotopes:

• Definition: Atoms of the same element that have the same atomic number (same number of protons) but different mass numbers (number of neutrons).

• Example: Cl-35 and Cl-37.

Problem Solving:

Calculate the average atomic mass of potassium, given the following abundances:93.1% occurs as K-39 and 6.9% occurs as K-41.Calculation: (39 x 0.931) + (41 x 0.069) = 39.1 amu.

• Definition of amu: atomic mass unit.

Radio-isotopes:

• An isotope that spontaneously decays to produce two or more smaller nuclei and radiation.

• Radioactive Decay: Spontaneous disintegration of unstable isotopes.

• Nuclear Radiation: Very small particles emitted during decay.

  • The 3 most common types:

    1. Alpha Particles: Helium atom.

    2. Beta Particles: Negatively charged particles identical to electrons.

    3. Gamma Rays: High-energy electromagnetic radiation.

Uses of Radio-isotopes:

1. Medical treatment (e.g., Iodine-131 for thyroid treatment).

2. Cancer treatment using gamma rays.

3. Dating fossils (e.g., Carbon-14).

Types of Radiation:

• Non-Ionizing Radiation: Moves atoms in a molecule, but does not remove electrons (e.g., radio waves, visible light).

• Ionizing Radiation: Has enough energy to knock electrons out of atoms, potentially damaging living tissue and DNA (e.g., x-rays).

The Periodic Table:

• Horizontal rows = Periods.

• Vertical columns = Groups.

• There are 8 main groups:

  • Group 1: Alkali Metals

  • Group 2: Alkaline Earth Metals

  • Group 7: Halogens

  • Group 8: Noble Gases

• At room temperature, only Bromine (Br) and Mercury (Hg) are liquids; all others are solids or gases.

• 5 types of gases are: HON + F, Cl

• Diatomic Molecules: H2, O2, Br2 F2, I2, N2, Cl2,

• Metals are on the LHS (left hand side) of the P.T., while non-metals are on the RHS (right hand side).

  Metals	Non-metals
  Conduct electricity	Do not conduct electricity except graphite (a form of carbon)
  Have high tendency to lose electrons.	Have high tendency to gain electrons.

  - Unique properties of transition metals (elements between groups 2 and 3):

  1. Very high melting point and boiling point

  2. They form more than one type of charge, for example iron can have +2 charge and +3 charge(Fe+2, Fe+3)

  3. They form colored compounds, for example

  Fe+2 forms green compounds , Cu+2 forms blue compounds

Valence Electrons:

• Electrons in the last shell of an atom.

• Example: Carbon has 4 valence electrons.

• Elements in the same group have similar chemical properties.

Periodic Trends in Atomic Properties:

1. Atomic Radius: Decreases across a period, increases down a group.

2. Ionization Energy: Increases across a period, decreases down a group.

3. Electron Affinity: Increases across a period, decreases down a group.

4. Electronegativity: Increases across a period, decreases down a group.

Example:

• For Na (190 pm) vs. Na+ (95 pm) and Cl (79 pm) vs. Cl- (181 pm), the radii differ due to ionization.

Conclusion:

• The periodic law states that the properties of elements change gradually as atomic numbers increase.

• Noble gases are stable and inert due to full electron shells.

• Lewis Symbol: Represents only the valence electrons of an element/atom.

Ch 2 Chemical Compounds and Bondings

Chemical bond: force of attraction between 2 atoms or 2 ions.

There are 2 main types of chemical bonds:

1. Ionic bond: a bond between a metallic atom (metal) and a non-metallic atom (non-metal) in which there is a complete transfer of electrons from the metal to the non-metal.

The compound which is formed is called an ionic compound.

Ex.: NaCl , MgCl2, Al2O3

To write the formula of an ionic compound we use the criss-cross method (we down cross multiply the charges without the sign, only the numbers of the charges)

Remark: if the charges are the same then they cancel each other in the formula so there will be one atom of the metal and one atom of the non-metal in the compound.

Exercise:

Write the formula of the compound which is formed between the following elements, and name each compound.

a) Rb and S:  

b) Ca and Se:

c) Al and Br:

d) Na and N:

- Draw a Bohr diagram to show the transfer of electrons (loss / gain) in an ionic compound.

Example: Na2O (sodium oxide)

Remark: The ionic bond is also described as an electrostatic force of attraction between a positive ion and a negative ion (eg: Na+           Cl- ).

2. Covalent bond: a bond between a non-metal and another non-metal in which there is a sharing of electrons between the non-metallic atoms. The compound that is formed is called a covalent or molecular compound.

Molecule: 2 or more atoms (non-metallic) bonded together; the atoms can be of the same element such as O2 or of different elements such as HCl, CH4, H2O,…..

We show the sharing of electrons between non-metallic atoms by using the Lewis diagram. In addition we can make intersecting circles for the atoms to show the sharing.

Reminder: Lewis diagram of an atom shows only the valence electrons of that atom.

Most of the atoms follow the octet rule (there are very few exceptions), that is each atom will have eight electrons in the valence shell (same as noble gases) except hydrogen will have 2 electrons after sharing (same as helium).

Remark: 2 electrons that are not bonded to any other atom is called a lone pair (non-bonding pair) of electrons.

Exercise:

Draw lewis diagrams to show the sharing of electrons in each of the following compounds:

1. NF3                                           2.  CH4                                   3. CO2

                 4. CCl4                            5. CH2O

Remark:In drawing Lewis structure, we show the bonds between the atoms and we also show all lone pairs (if present) on any atom.

- Naming molecular compounds:

1: mono                        2: di                        3: tri                     4: tetra                   5: penta                    6: hexa                         7: hepta                   8: octa                  9: nona                 10:deca

Example: PCl5 : phosphorus pentachloride

Remark: If the first element contains only one atom we don’t write mono before it; If the second element contains only one atom we have to write mono before it.

Ex.: NO : nitrogen monoxide            CO: carbon monoxide

Exercise: Fill in the table below

Remark: If the compound contains a transition metal, then we have to mention the type of charge of the transition metal by inserting a roman numeral in brackets (I, II, III, IV ….) after the symbol of the transition metal.

Example:

Name the following compounds:

- FeCl2 : Iron (II) chloride

- Cu(NO3)2 : Copper (II) nitrate

Remark: There are few transition metals that have only one type of charge such as zinc, nickel, and silver; in this case no roman numeral is required.

Zinc : Zn+2             Silver: Ag+                Nickel: Ni+2

- ZnSO4 : zinc sulfate

- AgNO3 : silver nitrate

- CuSO4 : copper (II) sulfate

- Co(NO3)3 : cobalt (III) nitrate

Exercise: Write the formula of the following compounds.

Calcium phosphate:                                   Iron(III) hydroxide:

Sodium hydroxide:                                   Manganese(II) hydroxide:

Barium sulfate:                                         Zinc carbonate:

Ammonium nitrate:

Remark: We must enclose the polyatomic ion in brackets if the number after it is more than 1.

Note: If the polyatomic ion that ends with the prefix –ate decreases by one oxygen atom then the prefix changes to -ite.

If the prefix ending with – ite decreases by one oxygen atom then we precede the prefix by hypo, whereas if the prefix ending with – ate increases by one oxygen atom then we precede the prefix by per.  

Example:

ClO3- is called chlorate; if we reduce one oxygen atom then the ion becomes ClO2- and is called chlorite, however if we increase by one oxygen atom then the ion becomes ClO4- and is called perchlorate;

and if the chlorite is reduced by one oxygen atom then the ion becomes ClO- and is called hypochlorite.

Exercise: Name the following compounds:

K2SO3 :

NaNO2 :

Mg(ClO4)2 :

LiBrO2 :

- Comparison Table between ionic and covalent (molecular) compounds:

 Chapter 3: Molecular Compounds and Intermolecular Forces.

- Many consumer products are made from molecular compounds.

- The raw materials (original material) used in the production of consumer products are classified as either renewable or non-renewable.

- Renewable resources is a resource that replenish itself as it is used (sustainable). Examples include: solar , water, wind,….

- Non-renewable resources is not replaced as quickly as it is consumed. Examples: crude oil (fuel)

- Most synthetic molecular compounds are produced from fossil fuels (petrochemicals).

- Recycling: converting a product back into material that can be used to make new goods.

- Biodegradable: capable of being broken down (decomposed) rapidly by the action of moisture, heat and micro-organisms (bacteria); for example plastics.

- Compostable: the ability of a material to decompose naturally , resulting in a product that is able to sustain plant life; for example: food scraps, cotton, wool, and wood.

Polar covalent and non-polar covalent bonds.

Non-polar covalent : pure covalent in which the electrons in the bond are shared equally (evenly) between the atoms.

These atoms are usually identical atoms but some non-identical atoms can have equal sharing of

electrons or close to equal sharing.

The equal sharing of electrons is due to the fact that the atoms have the same electronegativity or very close electronegativity.

Example: H – H, Cl – Cl, C – S,…..

Polar covalent : unequal or uneven sharing of electrons, electrons tend to shift (or move) more towards one atom than the other.

Example: H – Cl , H – O, N – H, H – Br, C – F, …..

Note: The difference in electronegativity between the atoms is what determines if the bond is polar or non-polar.

In summary, a polar bond is a covalent bond but with some ionic character.

- The atom with higher electronegativity will acquire a partial negative charge, whereas the atom with lower electronegativity will acquire a partial positive charge.

- These partial charges are called dipoles.

Example : H2O

Remark: Always the positive dipole is for the less electronegative element.

The negative dipole is for the more electronegative element.

Exercise (HW): Determine whether the covalent bonds, in each of the following molecules, are polar or non-polar.

a) CS2              b) CH4              c) HBr            d) SiO2      e) PCl3

Remark: If the bonds are non-polar covalent this implies that the molecule is non-polar, but if the bonds are polar then the molecule can either be polar or non-polar depending on the 3 dimensional shape of the molecule.

Intermolecular (IM) Forces.

IM Force: a force of attraction between molecules;

That is between an atom of one molecule (negative dipole) and an atom of another molecule (positive dipole). 

Remark: The bonds within a molecule (same molecule) is stronger than bonds between molecules.

Types of IM forces:

1. dipole-dipole force

This occurs between polar molecules, as in the above diagram.

2. A special type of dipole-dipole force called the hydrogen bond.

Hydrogen bond is a bond (bridge) that connects hydrogen to two atoms of the following elements: nitrogen, oxygen or fluorine (NOF); one atom in each molecule.

Exercise: Show the hydrogen bonding between two NH3 molecules

3. London dispersion forces (Van der Waal forces): 

These forces occur between non-polar (pure) covalent compounds (example: Cl2).

At a moment (instant) the distribution of electrons of the bond is uneven (unequal) between the atoms (example : 49.5 % - 50.5 %) and this produces temporary dipoles.

These dipoles are also described to be instantaneous or momentarily.

This type of IM force is the weakest type.

Order of increasing strength of IM forces  (weakest to strongest):

London dispersion , dipole-dipole forces, hydrogen bonding.

Unique (Special) Properties of Water.

- The following properties are due to the presence of hydrogen bonding between the water molecules.

1. High specific heat capacity:

It takes a lot of heat (energy) to raise the temperature of 1 gram of water by 1oC.

2. High surface tension.

The molecules of water tend to bond together after breaking (bonds reform between the molecules).

Example: striders walking on the surface of water.

Remark: This property is also known as cohesive forces (liquid-liquid).

3. Capillary action.

This is water travelling up narrow tubes (example from roots up the stem to the leaves).

4. Universal Solvent.

This is the fact that water has the ability to dissolve a wide range of substances.

5. High melting and boiling points.

This is in comparison with other molecular compounds.

Ex.: b.p. of H2O : 100 oC,

whereas b.p. of H2S is – 60 oC.

• Green Chemistry in Action.                

- The aim of green chemistry includes the following:

1. Sustainability . Using renewable energy resources rather non-renewable ones.

2. Safety. Producing or using less toxic chemicals (reduce toxicity).

3. Energy efficiency. Carrying out processes at lower temperatures or turning waste into useful energy.

4. End-of-life degradation. Designing products that degrade (break down) into harmless substances after use.

- Many products in use can be reused, recycled or biodegraded.

- Making products from recycled materials helps to reduce the quantity of garbage going to landfills.

- Green cleaning products (products that use less toxic chemicals) are increasingly available and effective.

Remark: A widely used biodegradable product is paper-based packaging.

- Scientists are trying to replace the normal flame resistant plastic with bioplastics which are way less toxic .

Few additives were to be added to bioplastics to make them less flammable.

Chapter 1: Atomic Structure and the Periodic Table Overview

• Atomic Structure:

  • Atomic Number: Number of protons in an atom.

  • Mass Number: Total number of protons and neutrons.

  • Subatomic Particles:

    • Protons: 1 amu, positive charge, inside nucleus.

    • Neutrons: 1 amu, no charge, inside nucleus.

    • Electrons: 0.0005 amu, negative charge, outside nucleus in shells.

• Ions:

  • Cation: Positive ion (loss of electrons).

  • Anion: Negative ion (gain of electrons).

• Isotopes: Atoms of the same element with different mass numbers (e.g., Cl-35 and Cl-37).

• Periodic Table Trends:

  • Atomic Radius: Decreases across periods, increases down groups.

  • Ionization Energy: Increases across periods, decreases down groups.

  • Electronegativity: Increases across periods, decreases down groups.

• Chemical Bonds:

  • Ionic Bond: Transfer of electrons between metals and non-metals (e.g., NaCl).

  • Covalent Bond: Sharing of electrons between non-metals (e.g., H2O).

• Unique Properties:

  • Transition metals have variable charges and form colored compounds.

  • Noble gases are inert due to full valence electron shells.