Chapter 3. Atoms: Building Blocks of Matter
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1 Chapter 3 Atoms: Building Blocks of Matter Atom: means, from Democritus (Greek, 400BC) Atom: smallest particle of an element that retains the of that element Chemical Reaction: transformation of substances into one or more Up to late 1700 s, info about reactions were Technology improved scales/balances so that analysis would be more accurate. This helped the discovery of some laws- 1
2 Law of Conservation of Mass: states that mass is neither nor during a chemical or physical reaction. i.e. total mass before = total mass after Law of Definite Proportions: states that a chemical compound contains the same elements in exactly the same, regardless of the size of the sample i.e. Water is H 2 O. Hydrogen comprises % of the total mass, and Oxygen makes up % of the total always! 2
3 Law of Multiple Proportions If two or more different compounds are composed of the same two elements, then the ratio of the masses is always a whole number CO 2 (carbon dioxide) vs. CO (carbon monoxide) CO 2 = 32g oxygen = 12g carbon CO = 16g oxygen = 12g carbon i.e. There is the oxygen in Carbon Dioxide than in Carbon Monoxide Dalton s Atomic Theory In 1808, John Dalton (English teacher guy) proposed an explanation for these laws. 3
4 1) All matter is composed of very small particles called 2) Atoms of the same are identical size, mass, etc Atoms of different are different. Duh!! 3) Atoms cannot be subdivided, created or 4) Atoms of different elements combine in simple to form chemical compounds 5) In chemical reactions, atoms are combined, separated or Dalton s Atomic theory helps explain the laws Law of Conservation of Mass Law of Definite Proportions 4
5 Modern Atomic Theory Dalton thought atom was & There are 3 major differences between Dalton s and the Modern theory 1) Atoms are made up of smaller particles protons, neutrons & 2) Atoms can be changed from one element into another by 3) Atoms of the same element are not all exactly the same - Structure of the Atom Atom: smallest particle of an element that retains the chemical properties of that element Atom consists of particles that occupy the small and the larger surrounding area at the centre of the atom contains (positively charged) and (negatively charged) occupy a region surrounding the nucleus 5
6 Discovering the Electron Late 1800 s, experiments were carried out on tubes Electric current passed through a gas at low Pressure producing J.J. Thompson (English guy) used this device to calculate the of these cathode ray particles to their mass, which he found to be negatively charged and constant These cathode ray particles are now known as 6
7 Since atoms are, atoms must contain charge as well. Thompson believed that the atom was a cloud of positive charge with negative electrons embedded into it. This is known as the Model. 1909: Robert Millikan (American physicist) showed that mass of electron ~ 1/2000 mass of Hydrogen atom (its actually 1/1837 th ) Millikan s experiments confirmed that electrons are, are present in all atoms, and that atoms are. 7
8 Millikan s Oil Drop Experiment Also: 1) Atoms are neutral, therefore, there must be a positive charge to balance the negative electrons 2) Electrons are a small % of the atom s mass, therefore there must be other particles that make up most of the atom s mass 8
9 The Nucleus In 1911, Ernest Rutherford ( a Kiwi) did an experiment where he bombarded a thin Gold foil with small, positively charged particles They believed the atom s mass & charge were evenly distributed Rutherford, along with Geiger & Marsden, expected the alpha particles to simply, like a small ball bearing through a thin veil, but some were actually back 9
10 Rutherford concluded that the atom consisted of a very small, dense centre that is charged, which he called the Also, that the nucleus was surrounded by a very large empty region in which the occupied, like planets around the sun Later, it was calculated that mass of a proton is 1836 times the mass of an electron, even though they have but charge 10
11 The neutron is just slightly heavier that a proton by 1/10 of a %, but is All atoms, except Hydrogen, contain An atom is neutral because they contain equal numbers of and i.e. It is solely the that determine what type of element the atom is Protons, which are positively charged, can stay close to each other in the nucleus because of nuclear forces Some Facts: Particle Symbol Charge Mass(kg) Relative Mass Electron e x Proton p x Neutron n o x Size of atom is measured in picometres (pm) = m 11
12 Counting Atoms Atomic Number (Z): of an element is the number of in the nucleus of each atom e.g. every Carbon atom has 6 (and also 6 ) The Periodic Table is arranged in order of increasing Atomic Number The Atomic Number identifies the element (look at your Periodic Table or on the next slide) 12
13 Isotopes: are atoms of the same element that have different. Because all atoms of the same element must have the same number of (and ), they can only differ in the number of, which do not change the atoms identity e.g. Hydrogen consists of just 1 proton & 1 electron and is also known as But there is another kind of Hydrogen that consists of 1 proton, 1 electron & 1 neutron, called which is only 0.015% of all Hydrogen There is still another type of Hydrogen which has 1 proton, 1 electron & 2 neutrons, called (The 3 Isotopes of Hydrogen:) % 0.015% trace = radioactive 100, ,000 = 12.33yrs 13
14 Mass Number (A): is the total number of & in the nucleus of an atom So, for Hydrogen: Name Atomic #(Z) Mass #(A) # protons # neutrons protons+neutrons Protium Deuterium Tritium So, # Neutrons = Nuclide: general term for any isotope of any element Isotope/Nuclide Symbol p + e - n o Protium Deuterium Tritium
15 Isotopes of Boron:. protons electrons neutrons Atomic # (Z) Mass # (A). Symbol Name Atomic Masses Mass (not weight) of atoms is very, very small, i.e. around kg Better to use a relative (or ratio) scale to compare mass of atoms to each other, but you need a standard. Choose Carbon, which has: 6 protons & 6 neutrons, 12 6 C, also known as It has been assigned a mass of exactly 12 atomic mass units (amu),so: 1 amu = 1/12 mass of Carbon-12 = x kg 15
16 Suppose you have: 10 pens weighing ½ oz each, and 40 pens weighing ¾ oz each. What is the average mass of each pen? 10 x ½ oz = 40 x ¾ oz = Total mass = Average mass = oz for each pen Or, another way: 10 / 50 = % & 40 / 50 = % ( % x ½ oz) + ( % x ¾ oz) = ( x ½ oz) + ( x ¾ oz) = + = oz 16
17 Average Atomic Mass: is the weighted average of the atomic masses of the naturally occurring isotopes of an element Example from book: Copper % of all Copper is Copper-63 ( 63 29Cu) which has a mass of amu % of all Copper is Copper-65 ( 65 29Cu) which has a mass of amu Calculate Average Atomic Mass ( % x amu) + ( % x amu) ( x amu) + ( x amu) = amu = amu (book likes 2DP) Summary: Copper -63 Isotope s Name Copper-65 Isotope s Symbol. # protons # electrons # neutrons. Atomic (Z) # Mass (A) # % % abundance 30.83% x amu atomic mass x amu Average (weighted) Atomic Mass. amu 17
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