What is the chemical formula for water? Draw the structure of water. Write down all the types of bonding that you know of.

•What is the chemical formula for water?

Draw the structure of water.Write down all the types of bonding that you know of.

What allows this drop of

water to hang there

without falling?

Surface Tension

Water ClingsHigh Surface Tension – water

molecules form hydrogen bonds with each other creating cohesion

A water strider can walk on the surface of a pond.

Capillary Action – water molecules from hydrogen bonds with other polar molecules creating adhesion

Capillary action causes water to creep up a narrow glass tube and paper.

Hydrogen Bonding in WaterHydrogen bonding causes water to absorb a

large amount of heat before

its temperature increases

appreciably and also causes

it to lose large amounts of

heat before its temperature

decreases significantly.

(Heat Capacity)

High heat capacity allows

organisms to maintain a

constant body temperature.

Hydrogen bonding causes

liquid water to absorb a large

amount of heat to become a

vapour. Many organisms

(including humans) dissipate body

heat by evaporation of water from

surfaces (skin-sweating; tongue-panting)

Chemical Context of Life

Matter (space & mass)

Element; compound The atom Atomic number (# of

protons); mass number (protons + neutrons)

Isotopes (different # of

neutrons); radioactive isotopes (nuclear decay)

Energy (ability to do work); energy levels (electron states of potential energy)

Chemical Bonding

Covalent Double covalent Nonpolar covalentPolar covalentIonicHydrogenvan der Waals

Covalent Bonding

Sharing pair of valence electrons

Number of electrons required to complete an atom’s valence shell determines how many bonds will form

Ex: Hydrogen & oxygen bonding in water; methane

Covalent bonding

Practice some examples of Covalent Bonding using Lewis structures.

What other properties do covalent bonds have?

Polar/nonpolar covalent bonds

Electronegativityattraction for electrons

Nonpolar covalent •electrons shared

equally •Ex: diatomic H and O

Polar covalent•one atom more electronegative than the other (charged)•Ex: water

Polar/nonpolar bondsRecall ►

Covalent bonds may be polar or nonpolar

Nonpolar covalent – electronegativity difference = 0

Polar covalent – electronegativity difference is greater than 0 but less than 1.7

Molecular polarity – dependant on both bond polarity and molecular shape

Symmetrical molecules whose bonds are all polar are nonpolar.

Asymmetrical molecules are nonpolar of all bonds are nonpolar, and they are polar if at least one bond is polar.

Ionic bonding

High electronegativity difference strips valence electrons away from another atom

Electron transfer creates ions (charged atoms)

Cation (positive ion); anion (negative ion)

Ex: Salts (sodium chloride)

Ionic bonds

Practice drawing ionic compounds using lewis structures.

What are some other properties of ionic bonds?

Hydrogen bonds

Hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom (oxygen or nitrogen)

van der Waals interactions

Weak interactions between molecules or parts of molecules that are brought about by localized change fluctuations

Due to the fact that electrons are constantly in motion and at any given instant, ever-changing “hot spots” of negative or positive charge may develop

Water Polar~ opposite ends, opposite charges Cohesion~ H+ bonds holding molecules

together Adhesion~ H+ bonds holding molecules to

another substance Surface tension~ measurement of the

difficulty to break or stretch the surface of a liquid

Specific heat~ amount of heat absorbed or lost to change temperature by 1oC

Heat of vaporization~ quantity of heat required to convert 1g from liquid to gas states

Density……….

Density

Less dense as solid than liquid

Due to hydrogen bonding

Crystalline lattice keeps molecules at a distance

Acid/Base & pH

Dissociation of water into a hydrogen ion and a hydroxide ion

Acid: increases the hydrogen concentration of a solution

Base: reduces the hydrogen ion concentration of a solution

pH: “power of hydrogen” Buffers: substances that

minimize H+ and OH- concentrations (accepts or donates H+ ions)

Two major parts of an atom

Nucleus (not to scale)

Electron Cloud

Three Major Sub-Atomic Particles

• Protons• Neutrons• Electrons

a single, relatively large particle with a

positive charge that isfound in the nucleus

PROTON (p+)

THE PROTON

p+

• Fat (heavy)

• Positive (charge)

• Doesn’t move (lazy)

a single, relatively large particle with a

neutral charge that isfound in the nucleus

NEUTRON (N°)

THE NEUTRON

N°

• Fat (heavy)

• Neutral (charge)

• Doesn’t move (lazy)

a single, very small particle with a

negative charge that isfound in a “cloud” around the nucleus

ELECTRON (e-)

THE ELECTRON• Skinny (very light)

• Negative (charge)

• Moves a lot (runs around)

e-

Review: Subatomic Particles

e-

N°

p+

Please complete the following table

Protons Neutrons Electrons

Where are they found?

Mass

Charge (attitude)

Nucleus Nucleus Electron Cloud

Heavy Heavy Very Light

Positive Neutral Negative

The total mass of all of the subatomic particles

in an atom (but really # of protons and

neutrons)

ATOMIC MASS #(A)

the number of protons in an atom

(assuming the atom is neutral, # of p+ = # of e-)

ATOMIC NUMBER (Z)

Example: Sodium

Na11

22.99

Atomic # = # of protons

Atomic Mass # = p+ & N°

Another Notation

Atomic # = # of protons

Atomic Mass # = p+ & N°

To calculate the number of neutrons, subtract the atomic number (smaller)

from the atomic mass number (larger)

A – Z = # of neutrons

Ex: How many neutrons does Sodium have?

Mass # - Atomic # = #N°(You may need to round the atomic #)

23 - 11 = 12 N°Na11

22.99

Atoms of the same element that differ in

charge.(They have the same # of p+, but different # of e-)

ION

Positive Ions(cations)

Negative Ions(anions)

• Na+ (lost 1 e-)

• Ca2+ (lost 2 e-)

• Al3+ (lost 3 e-)

• Pb4+ (lost 4 e-)

• H+ (lost 1 e-)

• Cl- (gain 1 e-)

• O2- (gain 2 e-)

• P3- (gain 3 e-)

• S2- (gain 2 e-)

• OH- (gain 1 e-)

If an atom GAINS electrons, its overall charge

becomes more negative.If it LOSES electrons, its

charge becomes more positive

Atoms of the same element that differ in

mass.(They have the same # of p+, but different # of N°)

ISOTOPE

Isotopes are CHEMICALLY the SAME as atoms, but

DIFFER PHYSICALLY because they have different masses.

A few examples of isotopes…

Complete the following table

Protons Neutrons Electrons

Na+

Br w/ mass 84

O2- with mass 13