Atoms, Molecules, Formulas, and Subatomic Particles

Transcription

1 Introduction to Chemistry Chapter 5 1 Atoms, Molecules, Formulas, and Subatomic Particles The Atom: The smallest particle of an element that can exist and still have the properties of the element building blocks for matter Dalton s atomic theory 1. Elements are made of tiny particles called atoms. 2. All atoms of a given element are identical. 3. The atoms of a given element are different from those of any other element. 4. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative numbers and types of atoms. 5. Atoms are indivisible in chemical processes. That is, atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way the atoms are grouped together. An atom is about m (an angstrom Å) in diameter. An atom has a mass about g Compounds: Substances which in pure form have the same composition. Same atoms Same ratio of those atoms (same mass ratio) Law of Multiple Proportions Same percent composition of the different elements E.g., Water, regardless of where it is observed, always has a ratio of 2 hydrogen and 1oxygen. The mass percentage of hydrogen is always 11.2 % and the mass percentage of oxygen is always 88.8 %. Molecules: A group of two or more atoms that function as a unit. Molecules exist alone Covalently bonded share electrons Diatomic molecules contain two atoms (same or different)

2 2 Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2 (Brinclhof) H H O O N N H F H Cl Homoatomic molecules have the same atoms bonded together C n (diamond, graghite, fullerenes) P 4 (phosphorus) Heteroatomic molecules have different atoms bonded together H 2 O C 6 H 12 O 6 CH 4 Ionic Compounds: (Salts) A group of two or more atoms that function as a unit. Exist in crystals Ionically bonded electrostatic attraction between ions Formula unit NaCl MgCl 2 KBr Chemical Formulas: 1. Each atom present is represented by its element symbol. 2. The number of each type of atom is indicated by a subscript written to the right of the element symbol. 3. When only one atom of a given type is present, the subscript 1 is not written. Molecular formulas: Molecules Formula unit : Ionic compounds

3 3 The Structure of the Atom In the late 1890s, a physicist in England named J.J. (Joseph John) Thomson demonstrated that the atoms of any element can be made to emit tiny negative particles. Cathode ray tube gas discharge tube vacuum tube with two metal electrodes anode: electrode attached to positive side of electrical power source cathode: electrode attached to negative side of electrical power source gas pumped out to have low pressure turn on power source, small amount of gas remaining glows get rid of almost all the gas and the glow from the gas stops greenish glow given off by the glass walls of tube indicated that the electricity was still moving from - to + electrodes. The canal rays were determined to have a negative charge and called electrons (after George Stoney s suggestion in 1891) Ernest Rutherford (gold foil experiment determined the positive charge was concentrated in the center of the atom. Nuclear atom an atom with a dense center of positive charge (the nucleus) around which tiny electrons moved in space that was mostly empty. Nucleus is made of proton Rutherford and a co-worker, James Chadwick, were able to show the 1932 that most nuclei also contain a neutral particle that they named the neutron.

4 4 The Fundamental Particles: 1 amu = x g The proton and neutron is 1836 times larger than the electron. The majority of the mass of an atom is located in the nucleus Atomic number: Z number of protons in the nucleus of an atom all atoms of the same element have the same number of protons Atomic mass number: A number of protons and neutrons in the nucleus of an atom Isotopes: atoms of the same element with the same number of protons but different numbers of neutrons isotopic symbol (nuclide) A X Z x = symbol of the element (protons + neutrons) (protons) = number of neutrons A - Z + n o 12C 13C 14C number of n o carbon 12 carbon 13 carbon 14 Neutral atoms (no charge) have the same number of electrons as protons Carbon has 6 protons and 6 electrons (if neutral)

5 Atomic mass: 5 relative mass Ex. Atom A is four times heavier than atom C. C is five times heavier than Atom B. What are the relative masses? Assume the mass is in amu. Assign atom B 1.0 amu. (The least weight) 1.0 amu x 5 = 5.0 amu for C 5.0 amu x 4 = 20.0 amu for A Carbon-12 has been the standard in which to determine the relative mass of atoms of other elements. 12 C has been given mass of amu. All other atoms are compared to the carbon-12 atom and are assigned amu according to their relative masses. Ex. Carbon is 12 times heavier than hydrogen but about 3 times lighter than argon. 1. If carbon is 12 amu then hydrogen must be 1/12 of 12, therefore, hydrogen must be 1 amu. 2. If carbon is 3 times lighter than argon, then argon must be 3x larger, therefore, argon must be 3x12= 36 amu Average mass (weights) of the elements. atoms of any one element may have several isotopes (same number of p + but different numbers of n 0 ). Therefore, a sample of an element cannot have an atomic mass just based on the mass of a single atom. The average atomic mass is bases on the percentage of the natural occurrence of the isotopes of an element. Ex. Oxygen has three isotopes which occur naturally in the following percentages: 16 O occurs at a % natural abundance 17 O occurs at a % natural abundance 18 O occurs at a % natural abundance For any sample of oxygen atoms, we can expect that the above percentages of isotopes will be represented. Therefore, the average of the atomic masses (weights) of each isotope is the accepted amu for oxygen. Determined by mass spec, the following amu where determined for each isotope. Isotope % Natural Abundance Mass (amu) 16 O O O The average atomic weight of a sample of oxygen would be: ( x ) + ( x ) + ( x ) = = amu

6 6 IONS: Examples: periodic table group I, II, III, V, VI,VII, and VIII Groups I III form cations Groups V VII form anions Group I form +1 charges, Group II form +2 charges, Group III form +3 charges CATIONS: (arrow means the atom has changed to an atom with one less electron) Na 0 Na + + 1e - sodium sodium ion K o K + + 1e - potassium potassium ion Mg o Mg emagnesium magnesium ion Ca o Ca e - calcium calcium ion Al o Al e - aluminum aluminum ion ANIONS: F o + 1e - F fluorine fluoride Cl o + 1e - Cl chlorine chloride O o + 2e - O 2 oxygen oxide N o + 3e - N 3 nitrogen nitride

7 Nuclear table: 7 Name of atom or ion Symbol of atom or ion Number of protons Number of electrons Number of neutrons Atomic mass Isotopic symbol beryllium 4 5 S 2 32 bromide 45 Cs Cs