Page | 1 MATS1101 Chemistry notes semester 2 2012 TOPIC 1 ...
MATS1101 Chemistry
MATS1101 Chemistry notes semester 2 2012 TOPIC 1: INTRODUCTION Atomic and molecular structure Elementary atomic structure The main points of Dalton’s Atomic Theory (1803), as accepted today, are: 1. 2. 3. 4. 5. 6.
every element is made up of atoms all atoms of any element are the same atoms of different elements are different (size, properties) atoms of different elements can combine to form compounds in chemical reactions, atoms are not made, destroyed or changed in any compound, the numbers and kinds of atoms remain the same
Using this theory we can explain three fundamental laws of chemical behaviour: 1. Law of Conservation of Mass and Energy: Matter is neither created or destroyed in a chemical reaction. Energy is neither created or destroyed in a chemical reaction, but it may be transformed from one form to another. In a chemical reaction the arrangement of the atoms is merely changed. 2. Law of Definite Composition: In chemical compounds, the various elements are always combined in definite proportions by number (and by mass). E.g. in water H2O, two hydrogen atoms are always combined with one oxygen atom. 3. Law of Multiple Proportions: Elements may combine together in more than one proportion in number or mass, but their compositions can always be expressed as a ratio of small integers. E.g. water H2O: H:O = 2:1 hydrogen peroxide H2O2: H:O = 1:1. Fundamental Sub-atomic Particles To understand chemical reactions we only need consider the three Fundamental Sub-atomic Particles: -
Electron→ has a mass of 9.1 x 10-28 g, and a charge of –1.602 x 10-19 C. Proton → has a mass of 1.67 x 10-24 g, and a charge of +1.602 x 10-19 C. Neutron →has a mass of 1.676 x 10-24 g, and a zero charge.
An atom is mainly free space, with a tiny, dense nucleus containing nearly all the mass (protons and neutrons) that is surrounded at some considerable distance by the electrons. Isoelectronic species have the same number and arrangement of electrons the Atomic Number of an element is equal to the number of protons in the nucleus of the atom The Mass Number is an integral number used to indicate the approximate mass of the atom = sum of protons and neutrons in the nucleus. Example: 126C is carbon, atomic number = 6, mass number = 12. (Also called carbon-12). It comprises 6 protons, 6 neutrons, and 6 electrons.
Page | 1
MATS1101 Chemistry
Isotopes -
Isotopes are atoms of an element with different numbers of neutrons. Thus they have the same number of protons, but have different mass numbers. Applications include: geological dating, isotopes for labelling in medicine and many other fields, therapy in medicine.
Neutral atoms: no. protons = no. electrons Ions: charged particle (atom or group of atoms)→when an atom gains or loses electrons it acquires a net charged, forms an ion Anions: negatively charged species (e.g. CL-, NO3-) no. protons < no. electrons Cations: Positively charged species (e.g. Na+, NH4+) no. protons > no. electrons
Atomic and molar mass -
A mole of a substance is the mass of it, in grams that contains the same number of particles as there are in exactly 12.000g of carbon-12. This is AVOGADRO’s number and is equal to 6.022 x 1023 particles A mole of a substance is its atomic, molecular or formula mass expressed in grams Amu = atomic mass unit, where 1amu= 1/12 mass of 1 atom of Carbon-12 (Unit also called the Dalton D) Atomic mass = sum of all isotopes (isotope mass [amu] x % abundance/100) Example: The composition of natural carbon is 98.892% carbon-12 (12.00000 a.m.u.) and 1.108% carbon-13 (13.00335 a.m.u.). Calculate the atomic mass of natural carbon. Atomic mass of natural carbon = (contribution of carbon-12) + (contribution of carbon-13) = (12.00000 x 98.892 / 100) + (13.00335 x 1.108 / 100) = 11.867 + 0.144 = 12.011 a.m.u.
-
molar mass of a molecule is the sum of the individual atomic masses. Molar mass (M) = sum of all elements (atomic mass (amu) x composition number) → units Da
Nomenclature and Stoichiometry -
compound: pure substance made from 2 or more chemically bonded atoms of different elements binary compound: contains 2 elements – Ax By molecule: stable entity of 2 or more chemically bonded atoms (may be of the same element)
Describing the composition by mass of HCN. If we took exactly 1 mole of HCN (27.0 g), then it would contain: H 1.0 g = 1.0 g / 27.0 g x100% H = 3.7 % H by mass C 12.0 g = 12.0 g / 27.0 g x 100% C = 44.4% C by mass N 14.0 g = 14.0 g / 27.0 g x 100% N = 51.9% N by mass.
Page | 2
MATS1101 Chemistry notes semester 2 2012 TOPIC 1: INTRODUCTION Atomic and molecular structure Elementary atomic structure The main points of Dalton’s Atomic Theory (1803), as accepted today, are: 1. 2. 3. 4. 5. 6.
every element is made up of atoms all atoms of any element are the same atoms of different elements are different (size, properties) atoms of different elements can combine to form compounds in chemical reactions, atoms are not made, destroyed or changed in any compound, the numbers and kinds of atoms remain the same
Using this theory we can explain three fundamental laws of chemical behaviour: 1. Law of Conservation of Mass and Energy: Matter is neither created or destroyed in a chemical reaction. Energy is neither created or destroyed in a chemical reaction, but it may be transformed from one form to another. In a chemical reaction the arrangement of the atoms is merely changed. 2. Law of Definite Composition: In chemical compounds, the various elements are always combined in definite proportions by number (and by mass). E.g. in water H2O, two hydrogen atoms are always combined with one oxygen atom. 3. Law of Multiple Proportions: Elements may combine together in more than one proportion in number or mass, but their compositions can always be expressed as a ratio of small integers. E.g. water H2O: H:O = 2:1 hydrogen peroxide H2O2: H:O = 1:1. Fundamental Sub-atomic Particles To understand chemical reactions we only need consider the three Fundamental Sub-atomic Particles: -
Electron→ has a mass of 9.1 x 10-28 g, and a charge of –1.602 x 10-19 C. Proton → has a mass of 1.67 x 10-24 g, and a charge of +1.602 x 10-19 C. Neutron →has a mass of 1.676 x 10-24 g, and a zero charge.
An atom is mainly free space, with a tiny, dense nucleus containing nearly all the mass (protons and neutrons) that is surrounded at some considerable distance by the electrons. Isoelectronic species have the same number and arrangement of electrons the Atomic Number of an element is equal to the number of protons in the nucleus of the atom The Mass Number is an integral number used to indicate the approximate mass of the atom = sum of protons and neutrons in the nucleus. Example: 126C is carbon, atomic number = 6, mass number = 12. (Also called carbon-12). It comprises 6 protons, 6 neutrons, and 6 electrons.
Page | 1
MATS1101 Chemistry
Isotopes -
Isotopes are atoms of an element with different numbers of neutrons. Thus they have the same number of protons, but have different mass numbers. Applications include: geological dating, isotopes for labelling in medicine and many other fields, therapy in medicine.
Neutral atoms: no. protons = no. electrons Ions: charged particle (atom or group of atoms)→when an atom gains or loses electrons it acquires a net charged, forms an ion Anions: negatively charged species (e.g. CL-, NO3-) no. protons < no. electrons Cations: Positively charged species (e.g. Na+, NH4+) no. protons > no. electrons
Atomic and molar mass -
A mole of a substance is the mass of it, in grams that contains the same number of particles as there are in exactly 12.000g of carbon-12. This is AVOGADRO’s number and is equal to 6.022 x 1023 particles A mole of a substance is its atomic, molecular or formula mass expressed in grams Amu = atomic mass unit, where 1amu= 1/12 mass of 1 atom of Carbon-12 (Unit also called the Dalton D) Atomic mass = sum of all isotopes (isotope mass [amu] x % abundance/100) Example: The composition of natural carbon is 98.892% carbon-12 (12.00000 a.m.u.) and 1.108% carbon-13 (13.00335 a.m.u.). Calculate the atomic mass of natural carbon. Atomic mass of natural carbon = (contribution of carbon-12) + (contribution of carbon-13) = (12.00000 x 98.892 / 100) + (13.00335 x 1.108 / 100) = 11.867 + 0.144 = 12.011 a.m.u.
-
molar mass of a molecule is the sum of the individual atomic masses. Molar mass (M) = sum of all elements (atomic mass (amu) x composition number) → units Da
Nomenclature and Stoichiometry -
compound: pure substance made from 2 or more chemically bonded atoms of different elements binary compound: contains 2 elements – Ax By molecule: stable entity of 2 or more chemically bonded atoms (may be of the same element)
Describing the composition by mass of HCN. If we took exactly 1 mole of HCN (27.0 g), then it would contain: H 1.0 g = 1.0 g / 27.0 g x100% H = 3.7 % H by mass C 12.0 g = 12.0 g / 27.0 g x 100% C = 44.4% C by mass N 14.0 g = 14.0 g / 27.0 g x 100% N = 51.9% N by mass.
Page | 2
