cheMYSTERY

Worked example:

Calcite contains 40.0% calcium, 12.0% carbon and 48.0% oxygen. What is the empirical formula?

There are three steps to working this out:

Divide each mass (this is the percentage in this case) by the molar mass of the element.

                   Ca :          C  :       O

          40.0/40.1      :   12.0/12.0   :  48.0/16.0

              0.998 :       1.00 :     3.00              mol

     2.  Divide by the smallest number ( the smallest number is 0.998 in this answer).

         0.998/0.998   :  1.00/0.998  :   3.00/0.998

     3.  Round the answers up. If you have an answer that ends in .5, you have to times   all the answers by 2.

The simplest whole-number ratio of atoms of each element present in a compound. The empirical formula is always the same as the actual formula for ionic compound but the empirical formula is not always the same as the molecular formula in molecular compounds. 

Examples:

Sodium chloride (table salt)

ionic formula: NaCl

empirical formula: NaCl

Methane

molecular formular: CH4

empirical formular: CH4

Butane*

molecular formular: C4H10

empirical formular: C2H5

Avogadro’s number (N) is the number of atoms in 12.01 grams of carbon. In chemistry, this is equivalent to a mole. 

Its numerical value is 6.02 x 10²³ particles.

The atomic mass of each element on the periodic table contain 6.02 x 10²³ atoms.

The volume occupied by one mole of softballs would be about the size of Earth. 

One mole of Olympic shot put balls has about the same mass as that of Earth. 

One mole of hydrogen atoms laid side by side would circle Earth about 1 million times.

Remember the mole (mol) is a unit of measure for an amount of a chemical substance. A mole is also Avogadro’s number of particles, which is 6.02 x 10²³ particles.

1 mol = 6.02 x 10²³ units

We can use the mole relationship to convert between the number of particles and the moles of a substance.

Molar mass is the mass of 1 mole of a substance in grams. 

For example, the atomic mass of carbon is 12.01 amu, so:

the molar mass (mass per mole) of carbon is 12.01 g/mol.

Since nitrogen occurs naturally as a diatomic, N2: 

the molar mass of nitrogen gas is 2 x 14.01 or 28.02 g/mol.

The molar mass of a substance is the sum of the molar masses of each element.

EXAMPLES:

The molar mass of H2O

Hydrogen - 2 x 1.01 g/mol = 2.02 g/mol

Oxygen - 16.00 g/mol

Total - 18.02 g/mol

The molar mass of Cu(NO2)2

Copper - 63.55 g/mol

Nitrogen -  2 x 14.01 g/mol = 28.02 g/mol

Oxygen - 4 x 16.00 g/mol = 64.00 g/mol

Total - 155.57 g/mol

Using the triangles below, we can find:

moles x 6.02 x 10²³ = number of particles.

number of particles ÷ 6.02 x 10²³ = moles.

moles x molar mass = mass.

mass ÷ moles = molar mass.

mass ÷ molar mass = moles.

Some calculations are two step problems. This is indicated by the absence of the word ‘mole’ in the question.

At Standard Temperature and Pressure or STP (which is 0° Celsius and 1 atmosphere) 1 mol of any gas occupies 22.4 L. The volume occupied by 1 mol of gas (22.4 L) is called the molar volume.

moles x 22.4 L/mol = volume

volume ÷ 22.4 L/mol = moles

The density of gases is much less than that of liquids. The density of any gas at STP can be easily calculated, with this formula:

There are three interpretations for the mole:

1 mol = 6.02 x 10²³ particles

1 mol = molar mass

1 mol = 22.4 L at STP

This gives us three ways to convert between moles, particles, mass and volume.

The percent composition of a compound lists the mass percent of each element.

Electronegativity: the power of an atom to attract the electrons (electron density) in a covalent bond towards itself

Atoms are held together by forces of attraction between opposite charges. In a covalent bond, the electrons shared between the atoms will be unevenly spread if one of the atoms attracts electrons more strongly than the other.

Electronegativity is measured on the Pauling scale, which runs from zero (no electronegativity) to four (the most electronegative).

Some common values to learn are:

H – 2.1

C – 2.5

Br – 2.8

Cl – 3.0

F – 4.0

Electronegativity depends on the charge and the atoms;

The higher the nuclear charge, the more attraction between the nuclear charge and the pair of electrons in the covalent bond, so the more electronegative the atom.

The larger the atom, the less electronegative the atom – because the covalent bond is further from the nucleus so there is less attraction

The more shells in an atom, the less electronegative the atom – because there is more shielding between the nucleus and the covalent bond.

Trends in Electronegativity

Electronegativity increases going across a period – because the nuclear charge increases, the number of inner electrons remains the same and the atoms become smaller

Electronegativity decreases going down a group – because there is more shielding by electrons in inner shells and the atoms get larger (this has a greater effect on the electronegativity than the increase in nuclear charge).

The most electronegative atoms are in the top right-hand corner of the periodic table (but not the Noble Gases; they have zero electronegativity).

N, O and F are the most electronegative elements

H is one of the most electropositive elements

C and H have very similar electronegativity.

Bond Polarity

When both atoms in a covalent bond have the same electronegativity, the electron pair is shared equally and the bond is non-polar. (Usually, the two atoms have to be the same element for this to occur.)

When the atoms in a covalent bond have different electronegativity values, the more electronegative atom has a greater attraction to the electron pair. This means that the electron pair is pulled towards the more electronegative atom, which leads to a polar bond with an unequal sharing of the electron pair. This is shown using the symbols δ+ and δ-. (This is not a full ion.)

When an electron pair is shared unequally between two atoms, each atom gains a partial charge, shown by a lower-case Greek letter delta (δ). The more electronegative atom will gain a δ- charge and the less electronegative atom will gain a δ+ charge.

The polarity of the covalent bond depends on the difference in electronegativity;

Atoms with a large difference in electronegativity are ionic

Atoms with a small difference in electronegativity are polar covalent

Atoms with no difference in electronegativity are pure covalent.

Electronegativity - The measurement of an atoms attraction for shared pair of electrons in a bond

Across the period - Electronegativity increases as there is an increase in protons which gives a greater attraction for bonding electrons

Hybridization can be a difficult concept to master. There are so many different rules to follow that it can be hard to keep them straight. We will walk through an example with CH4 to illustrate a good way of going through hybridization problems.

1.      The best way to start is by drawing out the electron ground configuration of the atom. This shows just the valence electrons of each atom because those are the only electrons involved in forming bonds.

2.   This carbon seems like it should only be able to form two covalent bonds, so we make a new type of orbital, known as a hybrid orbital to allow Carbon to form four bonds. The 2s and 3 2p orbitals are combined to make four sp3 hybrid orbitals shown below. The orbitals are called sp3 because they are made up of 1 s orbital and 3 p orbitals, which explains the 3.

3.   Now it is easy to draw a picture of the orbitals and show the hydrogen atoms bonding to the carbon atom. Make sure to draw in each electron so that it is easy to see which electrons are shared in the covalent bond.

All images sourced: Chem Wiki

Here is a brief overview of common hybrid orbitals:

• Linear - sp
- the hybridization of one s
and one p
orbital generate two hybrid orbitals oriented 180°

apart.

• Trigonal planar – sp2

- the hybridization of one s
and two p
orbitals generate three hybrid orbitals oriented 120°

from each other all in the same plane.

Biology and chemistry are intertwined at intimate levels, so in order to make sense of what’s happening as we progress, you’re going to need some basic chemistry info. In case you have no chem background or you’re just really forgetful, I’m writing a couple of recap articles on necessary concepts.

Covalent bonds:

All atoms want to have eight electrons in their outermost shell—i.e., have eight valence electrons. To do this, two atoms sometimes “share” electrons to create the illusion that they both have eight. This is called a covalent bond. There are two different types, determined by the electronegativity of an atom: how much it wants to hog electrons.

1. Polar covalent bonds are formed when one atom is more electronegative than the other, so it pulls on the electron more strongly. This causes it to gain a partial negative charge, while the less electronegative atoms gain a partial positive charge. Water is a good example of this—see how the electronegativity “distorts” the bonds at an angle.

(Image Source)

2. Non-polar covalent bonds are formed when the electronegativity of the atoms are equal, so their shared electrons are shared equally. In this case, neither atom has an induced charge.

Ionic bonds:

These are formed when atoms the electronegativity of one atom is much stronger than the electronegativity of the other atom, so it’s not content with sharing—it just strips electrons from the other atom. The atoms then become ions: atoms that have lost or gained electrons. One is negatively charged (anion) and the other is positively charged (cation), and because they’re oppositely charged, they’re electrostatically attracted to each other and ionic bonds are formed. Note that the electron transfer isn’t the cause of the bond, but it does create the charge that the bond is based on.

Hydrogen bonds:

These intermolecular attractions are weaker than the bonds described, created when a hydrogen atom that is covalently bonded to one electronegative atom is attracted to another electronegative atom. Essentially, it’s an attraction between hydrogen (which, remember, is slightly positive) and an electronegative atom (which is slightly negative because it shares “more” of the electrons).

(Image Source)

It’s important to understand these bonds because they help explain how many biological molecules act. For example, water is polar and covalent, and this configuration is incredibly vital for life. When water freezes, the hydrogen bonds between its molecules create a rigid, lattice-like structure that separates out the atoms, thus making its solid form less dense than its liquid form. This is the reason that ice floats, and without this particular quirk of chemistry, it would be tough for marine organisms to survive. Usually, the ice that freezes first in cold conditions rises to the tops of lakes or oceans and forms an insulating layer, preventing the rest of the water from freezing and allowing organisms to survive the cold—if entire oceans froze every winter, life may have never evolved.

atoms sometimes take electrons from their bonding partners (ex: transfer of electron from sodium to chlorine)

After the transfer of an electron, both atoms have charges  

ion: charged atom 

cation: positively charged ion

anion: negatively charged ion

ionic bond: attraction between an anion and a cation