Chapter 4: Atoms and Elements

Transcription

1 C h e m i s t r y 1 2 C h 4 : A t o m s a n d E l e m e n t s P a g e 1 Chapter 4: Atoms and Elements Remember to keep up with MasteringChemistry, Workshops, Mini-Reports and Labs Early Ideas on Matter: Philosophers (Chinese- yin/yang; Greeks-earth/wind/fire/water) speculated about the nature of stuff without relying on scientific evidence Leucippus (fifth century BC) and his student Democritus ( BC) first suggested the material world when broken down to the extreme would consist of tiny particles called atomos, meaning indivisible. Alchemists through the middle ages physically experimented with matter aiming to create gold from base metals and an elixir for everlasting life. Englishman Robert Boyle ( ) is generally credited as the first to study the separate science we call chemistry and the first to perform rigorous experiments. Antoine Lavoisier ( ) discovered the mass of combustion products exactly equals the mass of the starting reactants. Law of Mass Conservation (Law of Conservation of Matter); Mass is neither created nor destroyed in chemical reactions Joseph Proust ( ) studied copper carbonate, the two tin oxides, and the two iron sulfides. He made artificial copper carbonate and compared it to natural copper carbonate, showing that each had the same proportion of weights between the three elements involved (Cu, C, O). He showed that no intermediate indeterminate compounds exist between the two tin oxides or the two iron sulfides. Law of Definite Proportions (Law of Constant Composition); Elements combine together in specific proportions. All samples of a given compound, regardless of their source or how they were prepared have the same proportions of their constituent elements. These early ideas led to the foundation steps in atomic theory. Atomic theories explain the behavior of atoms. We will cover Dalton s Indivisible atom, J.J. Thomson s Plum Pudding model, Rutherford s Nuclear model of the atom, the Bohr s Quantum (orbit) model that mathematically only works for one electron systems and the Orbital Wave Mechanical model. The first three models are found in Chapter 4 while the last two are found in Chapters 9.

2 C h e m i s t r y 1 2 C h 4 : A t o m s a n d E l e m e n t s P a g e 2 Dalton s Atomic Theory (1808): 1. Elements are composed of tiny, indivisible particles called atoms. 2. Atoms of a given element are identical in properties, but atoms of one element are different from the atoms of all other elements. 3. Compounds form when atoms of two or more different elements combine in whole number ratios. Chemical reactions do not create or destroy atoms, they are just rearranged. Dalton s atomic theory led to another scientific law Law of Multiple Proportions: When two elements form two different compounds, the masses of element (B) that combine with 1g of element (A) can be expressed as a ratio of small whole numbers. Example: CO(1 g C to 1.33 g O) vs CO 2 (1 g C to 2.67 g O) J. J. Thomson ( ); By the mid-1800 s new experiments gave data that was inconsistent with an indivisible atom. Cathode ray tubes (CRT) contain very low pressures of a gas and have high voltage passed through electrodes on either end. Experiments with CRT gave radiation that is negatively charged. The same negative charged substance that fluoresced (gave off light) was found using many different gases. By 1897, JJ Thomson published a paper that concluded the cathode rays are streams of negatively charged particles, later known as electrons. These particles are smaller than the atom itself, therefore these are the first sub-atomic particles identified through experiment.

3 C h e m i s t r y 1 2 C h 4 : A t o m s a n d E l e m e n t s P a g e 3 Plum Pudding Model; This experiment led to a divisible neutral atom which must have both negative and positive charges. JJ Thomson called his atomic theory the Plum Pudding Model of the atom. A positive sphere like pudding contains particles (plums) of negatively charged electrons. Since the atom is neutral, there must be a positively charged electric field as well. Thomson assumed there were no positively charged particles since none showed up in the experiment. He incorrectly predicted much of the mass of the atom comes from the mass of electrons. In 1909 Robert Millikin; Robert Millikin measured the charge of an electron ( x Coulombs) through an oil drop experiment performed numerous times over 5 tedious years. Using Thomson s charge to mass ratio ( x 10 8 C/g) the electron mass was accepted as x g, about 2000 times smaller than a single H atom. This caused the question: What is the major contributor of an atom s mass.

4 C h e m i s t r y 1 2 C h 4 : A t o m s a n d E l e m e n t s P a g e 4 Ernest Rutherford ( ): In 1910 Ernest Rutherford created the gold foil experiment experiment to test Thomson s Plum Pudding model. The results showed most of the heavy positive alpha particles passed right through a thin gold foil. Surprisingly, a small portion of alpha particles were deflected or even sent back. If Thomson s Plum Pudding atomic model was correct, this would be similar to a rifle shot through tissue paper, and no bullet should be deflected. Rutherford s Nuclear Model of the atom explains why some of the alpha particles were deflected of bounced back as the picture shows. The nuclear model has all the positive charge (protons) densely set in the center (nucleus) and the particles of electrons spread out in a cloud around the nucleus. 1. Most of the atom s mass and all of its positive charge are contained in a small core called the nucleus. 2. Most of the volume of the atom is empty space through which the tiny, negatively charged electrons are dispersed. 3. The number of negatively charged electrons outside the nucleus is equal to the number of positively charged particles (protons) inside the nucleus, so that the atom is electrically neutral. The dense nucleus makes up more than 99.9% of the mass of the atom, but occupies only a small fraction of its volume. The low mass electrons are distributed through a much larger region. A single grain of sand composed of just atomic nuclei would have a mass of 5 million kg. Astronomers believe that black holes and neutron stars are composed of this kind of incredibly dense matter.

5 C h e m i s t r y 1 2 C h 4 : A t o m s a n d E l e m e n t s P a g e 5 Neutrons: It did not make sense to have all the positive particles (protons) so close together in a nucleus, they would repel each other. Additionally, some mass was missing. One of Rutherford s students, James Chadwick ( ), proposed there are neutrons, neutral particles within the nucleus similar to protons. Neutrons were isolated later in Atomic Structure: What we have so far Particle Charge Mass (amu) Mass (g) Electron amu x g Proton amu x g Neutron amu x g 1 amu = x g Solve for the inverse of this number: amu = 1 g Atoms are extremely tiny with diameters around m: The tiny atomic nucleus is surrounded by a very large cloud of electrons The nucleus contains almost all the mass of an atom. It is positively charged and contains protons (+1 ) and neutrons (0 charge) Size Example: a marble (nucleus) in the center surrounded by a large football stadium (a cloud of electrons). Neutral atoms have the same number of electrons and protons. The number of protons defines the element. Each chemical element (X) has a unique number of protons (atomic number, Z). Ions have more or less electrons than protons. Cations lose electrons, are positive (metals) Anions gain electrons, are negative (nonmetals) Isotopes will be the same element with the same number of protons, but the number of neutrons are different. Isotopes are chemically identical. The protons plus neutrons is the Mass Number (A) Nuclide symbols (or Isotope Symbol Notation) indicate particular isotopes and ions. A Z X

6 C h e m i s t r y 1 2 C h 4 : A t o m s a n d E l e m e n t s P a g e 6 Example 1: Isotope Symbols Fill in the nuclide symbols chart. Nuclide symbol, name A Z X protons neutrons electrons atomic mass 12 6 C carbon carbon N Sulfide ion Potassium ion Elements: Many of our element symbols are based on its English name Some element names are based on Greek or Latin origins sodium, Na is from the Latin word natrium lead, Pb is from the Latin word plumbum which means heavy phosphorus, P is from the Greek words phôs (light) and phoros (bearer), Phosphoros was a god of light in Greek myth. Some element names honor locations or people Berkelium, Bk, for the location of the lab that created it first. Einsteinium, Es, for Albert Einstein Periodic Table: Patterns and the Periodic Law Development: 1869 Dmitri Mendeleev (Russia) and Lothar Meyer (Germany) classified known elements (about 65 known at that time) and noted similar physical and chemical properties were found periodically when arranged by increasing atomic weight and grouped together by chemical reactivity. Several holes led to predictions of elements and their properties that were not yet discovered eka-aluminum (Ga) and eka-silicon (Ge). Periodic Law when the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically

7 C h e m i s t r y 1 2 C h 4 : A t o m s a n d E l e m e n t s P a g e 7 The first periodic table: Ordered elements by atomic mass Put elements with similar properties in the same columns Used pattern to predict properties of undiscovered elements Where atomic mass order did not fit other properties, Mendeleev reordered by other properties Example: Te & I 1915 Henry Moseley developed the concept of atomic numbers. He improved the periodic table by ordering the elements by increasing atomic number. More holes were found, which led to the discovery of more elements and the family of noble gases. The periodic table gives us a great amount of information in an organized manner. Vertical columns are called groups or families. If you are aware of the properties of a couple elements in a group, you can make a good guess at the properties of the other elements in the same group. Periods are the horizontal rows in the periodic table. Many patterns can be seen or predicted following periods and groups. It is easy to identify certain expected characteristics by locations, such as Metals, Nonmetals Metalloids. Periodic Table: Organization: family/group period metals nonmetals metalloids/semiconductors, Groups: Main Group Transition Metals Inner Transition Metals or Actinides and Lanthanides Alkali Metals Alkaline Earth Metals Halogens Noble Gas Coinage metals

8 C h e m i s t r y 1 2 C h 4 : A t o m s a n d E l e m e n t s P a g e 8 Atoms may lose or gain electrons to form Ions: Cations (positive charge): Na + Anions (negative charge): Cl - Naming: Preview of Chapter 5: Cations with known oxidation state of metal Group 1A (+1), 2A (+2), Al and Ga (+3), Zn and Cd (+2), Ag (+1) Name of ion is identical to the name of the atom for cations Variable oxidation state of metal Transition metals and metals below the nonmetal on the right have a variable oxidation state that must be indicated by Roman Numerals in parenthesis (this method is what I expect you to learn. Fe +3, iron (III); Fe +2, iron (II); Cu +1, copper (I); Sn +4, tin (IV) An alternative method differentiates from the higher oxidation number and lower oxidation number using the old form of the name and ic or ous as an ending respectively. (you are not responsible for knowing the ic and ous ending of metal cations) Fe +3, ferric Fe +2, ferrous; Cu +2, cupric; Cu +1, cuprous; Sn +4, stannic; Sn +2, stannous Anions Naming: Group VA (-3); VIA (-2), VIIA (-1) Name of the element root followed by ide. N -3, nitride; S -2, sulfide, Br -1, bromide Atomic Weights: The atomic mass scale is arbitrarily defined by international agreement and is based a standard isotope carbon-12, defining its mass to be exactly 12 amu. Weighted average atomic masses take into consideration the natural abundance of all the isotopes of an atom. Masses and isotopic abundances are measured by Mass Spectroscopy.

9 C h e m i s t r y 1 2 C h 4 : A t o m s a n d E l e m e n t s P a g e 9 Mass Spectrum quantifies the results The mass spectrum for zirconium Isotopes: The 5 peaks in the mass spectrum shows that there are 5 isotopes of zirconium - with relative isotopic masses of 90, 91, 92, 94 and 96 on the 12 C scale. The abundance of the isotopes In this case, the 5 isotopes (with their relative percentage abundances) are: zirconium zirconium zirconium zirconium zirconium (This simple example rounds off the mass much more than I generally accept.)

10 C h e m i s t r y 1 2 C h 4 : A t o m s a n d E l e m e n t s P a g e 10 Working out the relative atomic mass Using the equation Weighted Atomic Mass = (0.515 x 90)+(0.112 x 91)+(0.171 x 92)+(0.174 x 94)+(0.028 x 96) = 91.3 is the relative atomic mass of zirconium. Mass: Simple vs. Weighted Average: Simple average: add all the numbers and divide by the count Solve for the simple average... Given: 12.0 g, 16.0 g, 17.0 g Weighted average: Mass = individual mass value x fractional abundance Weighted average takes into consideration the fractional abundance of each number. Fractional abundance is the decimal form of the percent abundance. All fractional abundance values add up to a total of one (1.00) so there is no reason to divide by the count. Solve for the weighted average Given: 12.0 g (80.0%), 16.0 g (15.0%), 17.0 g (5.0 %) Naturally occurring weighted masses for elements are found on the periodic table: Atomic mass = isotopic mass x fractional abundance Atomic mass = (mass A x fract. abund. A ) + (mass B x fract. abund. B ) + (. All the fractional abundance values add up to a total of one (1.00)

11 C h e m i s t r y 1 2 C h 4 : A t o m s a n d E l e m e n t s P a g e 11 Example 2: Use the weighted average to solve the average atomic mass found in nature for Si Given the following information on its naturally occurring isotopes Keep appropriate significant figures. 28 Si: amu 92.21% 29 Si: amu 4.70% 30 Si: amu 3.09% Example 3: There are two naturally occurring isotopes of chlorine. Calculate the percent abundance of each isotope given the following information on the masses and given that the naturally occurring weighted atomic mass of chlorine is amu 35 Cl: amu (1-x) 37 Cl: amu (x) Preview: Counting Atoms by Moles: Avogadro s number: x particles = 1 mole Converting atoms to moles Converting moles to atoms Molar Mass: Solving for molar mass of molecules and compounds O 2 H 2 O CoBr 3