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Basic Atomic Structure – History of the Atom Early atomic structure Greeks (400 – 300 B.C.) Continuous theory of matter – matter could be divided forever without reaching a single smallest indivisible unit The discontinuous theory of matter – there existed some smallet piece of matter John Dalton John Dalton (1766 – 1844) “Father of Chemical Atomic Theory” Dalton was the first individual to really expound upon the concept of the atom – this was a result of evaluating the work of other scientists as well as some of his own work involving the study of gases Dalton’s theory published over the years 1803 – 1807 consisted of several main ideas: Dalton’s Atomic Theory 1. Chemical elements are made up of atoms – Elements are made up of tiny, discrete, indivisible and indestructible particles called atoms. • Derived from the Greeks – These particles maintain their identity through chemical and physical changes. They are only rearranged. • Explained - Law of conservation of mass – developed by A. Lavoisier Note: Element – a chemical substance which cannot be further simplified by chemical means – a definition developed by A. Lavoisier Dalton’s Atomic Theory cont’d 2. The atoms of an element are identical - Atoms of the same element have the same properties, i.e. weight – Today, this concept has been proven incorrect with the knowledge that elements have isotopes. 3. Atoms of different elements have different properties - Atoms of different elements have different properties including different weights. – Again, this originated with the Greeks, however, it was Dalton that determined what some the weight values were. Dalton’s Atomic Theory cont’d 4. Atoms combine in small whole number ratios e.g. 1 to 1, 1 to 2, 2 to 3, etc. – This coincides with Joseph Proust’s Law of Definite Proportions which is also referred to as the Law of Constant Composition. – The law states that; compounds are composed of elements combined in a fixed ratio by mass. Dalton – other work Dalton also developed another law known as the Law of Multiple Proportions which states; that the ratio of a fixed mass of one element when combined with the multiple masses of another element may be expressed in small whole numbers. In simpler terms, the same element can unite in more than one way with another element to form several compounds. Dalton’s Model John Dalton’s model of the atom appeared as a spherical, indivisible and indestructible unit (~1807) Electricity Static electricity – is the build up of electrical charge on an object (recognize most objects are electrically neutral) - being static implies that it is not moving These objects of one charge are attracted to objects of the opposite charge, however, objects with the same charge repel each other This relationship between charged objects is described by the Law of Electrostatic Forces Electrostatic force is the force observed between two charged objects Electricity cont’d Electrically charged objects will attempt to return to their neutral condition. Why? This flow of electrical charge is referred to as electrical current. Electrical current is measured in units of amperes The ampere is defined as an electrical current which transports one coulomb of charge past a point in one second The unit coulomb, C, represents a quantity of charges equal to 6.24196 x 1018 charges Electricity cont’d The behavior of electrical charges is often observed using electrodes. Electrode – site or terminal of electrical charge Two types of electrodes; anode which is positive and cathode which is negative Coulomb’s Law Fel = kQ1Q2 r2 where Fel is electrostatic force; k is a constant with a value of 8.99 x 109Nm2/C2; Q1 and Q2 are the quantity of charge on the two objects; and r is the distance between the charged objects Understanding Coulomb’s Law Implications: - force is inversely proportion to the square of the distance between the charges - force is proportional to the product of the magnitudes of the charges - force is attractive for charges of opposite signs and repulsive for charges of the same sign Particles and Fields A charged particle moving through a field (electrical or magnetic) may have its path altered by the field – It will be attracted toward the opposite pole and repelled from the like pole The particle may only be deflected if the field is sufficiently weak, however, a strong field may actually result in capture of the particle It should be noted that deflection by a magnetic field is at right angles (perpendicular) to the field while the deflection by an electrical field is in the plane of that field Deflection of Particles Deflection of Nuclear Particles (note charges) Voltage The strength of a field or current is related to the potential energy of the electrons or quantity of charge(s) present This potential difference is measured in units of volts – it is the “push” behind electron movement The volt is defined as a potential difference of 1 joule per coulomb What is a joule? Voltage The greater the voltage, the greater the potential energy of the system, and the more work that may be done by that system Most household voltages are? Techniques for working with small moving charged objects Experiment frequently require evacuated or partially evacuated containers (metal or glass) as gaseous molecules present may interfere with the particle and its motion Particle observation usually requires the interaction/collision of the particle with another substance – use of a fluor or use of photographic film Cathode Ray Tube (CRT) OBSERVATIONS of the CRT If an object is placed in the cathode ray a shadow will be cast upon the wall at the glowing end of the tube – Straight line of travel CRT The cathode ray can push a small paddle wheel up an incline against the force of gravity – The ray carried energy and could do work The ray is deflected from straight line path by a magnetic field Possible relationship between the two as the relationship between magnetism and electricity hasn’t been discovered yet CRT Thomson Discovers the Electron J.J. Thomson deflects the cathode ray with an electrical field – The rays bend toward the positive pole, therefore, they are negative – 1897 announces that the corpuscles (electron) is a negatively charged particle as it is deflected towards the positive plate in an electrical field J.J. Thomson also Determined Using a modified CRT Thomson worked on determining the charge to mass ratio (e/m) of the electron The value obtained for the electron was: e/m = 1.759 x 108 C/g Charge to Mass Determination Robert Millikan Millikan performed the “oil drop” experiment in 1909 to determine the charge on the electron Millikan Oil Drop Millikan Oil Drop **Millikan’s results showed that the value of ed varied, but was always a multiple of a small number. - this resulted from the variation in ionizations **the small numerical multiple was 1.602 x 10-19 C **Millikan believed this to be the smallest possible charge, that belonging to the electron Mass of the Electron e___ e/m = = 1.602 x 10-19 C x m = 1.602 x 10-19 C 1.759 x 108 C/g 1 g______ 1.759 x 108 C 9.107 x 10-28 g Actual Electron Mass Actual mass of the electron is 9.109 534 x 10-28 g Memorize mass as 9.11 x 10-28 g Proton In 1886, Eugene Goldstein observed luminous rays streaming through a cathode whose modification was that it had holes It also contained hydrogen gas at low pressure The magnitude and direction of the rays in the presence of a magnetic field indicated they were positive, had a mass significantly larger than the electron, and travel at a much lower velocity Canal Rays Proton Credit for the discovery of the proton is debatable pending upon source - However, some people credit Rutherford with its discovery The proton’s mass was eventually identified as 1.672 65 x 10-24 g X-rays Discovered Wilhelm C. Roentgen (1845 – 1923) ~1895 - Accidentally discovered x-rays when his wife passed her hand in the path of x-rays revealing her bone structure and wedding ring This discovery inspired Becquerel to study fluorescence and phosphorescence The “Hand” Radioactivity Discovered A. Becquerel accidentally discovered radioactivity in 1896 while working on fluorescence and phosphorescence Worked with potassium uranyl sulfate – Radioactivity is defined as the spontaneous emission of subatomic particles or energy by disintegration of atomic nuclei The Ore Some types of radiation Alpha (α) radiation – composed of helium nuclei traveling at fairly high speeds – These particles carry a charge of +2 and have a mass of ~6.64 x 10-24 g – Alpha is a somewhat poor penetrating radiation – Equation symbol is 42He – Strong ionizing potential Beta (β) radiation – composed of high speed electrons – These particles carry a charge of –1 and have a mass of 9.11 x 10-28 g – Beta has fair penetrating ability – Equation symbol is 0-1e – Good ionizing potential Radiations cont’d X-rays – are composed of electromagnetic radiation – – They have good penetrating ability Poor ionizing potential Gamma rays – are similar to x-rays (electromagnetic radiation) but usually carry more energy – – – They have very strong penetrating ability Symbol is (γ) Variable ionizing potential Radiations Gold Foil Experiment Rutherford (1871 – 1973) along with coworkers Geiger, Marsden, and Bohr devised and performed the gold foil experiment in 1911 The experiment: A thin sheet of gold foil was enclosed by a fluor covered circular shroud – There was an opening in the shroud through which alpha particles were shot at the gold foil Gold Foil Experiment Results of Gold Foil Experiment 1) The vast majority of alpha particles passed through the foil without any deflection from their path 2) Some of the alpha particles passed through the foil with a minor deflection to their path 3) A scant few (1 in 8,000) alpha particles were deflected at large angles, in essence bounced backwards Rutherford’s Conclusions 1) Most of the alpha particles passed through the foil without deflection because the vast majority of the atom is empty space with a few electrons in it 2) The few alpha particles deflected at minor angles came close to a small bundle or core of positive charge which were, therefore, repelled 3) The scant number of alpha particles being bounced back was a result of an almost direct collision with a very massive (extremely dense) and positively charge region occupying a very small volume Atomic Models • Dalton – spherical indivisible unit • Atomic Models • Thomson – the plum pudding model • a “sea” of positive charge with electrons scattered throughout • a divisible atom – subatomic particles existed • an electrical nature associated with the atom Atomic Models Rutherford – planetary model - the atom contained a small, positively charged, dense core or center called the nucleus - the electrons traveled around outside the nucleus - most of the atom’s volume was actually empty space - the nuclear diameter is about 10-4 the diameter of the atom Problems w/ planetary model • Criticisms of the planetary model of the atom • Why weren’t the electrons pulled into the positively charged nucleus of the atom? – Rutherford responded by stating the electron’s motion prevents it from being pulled into the nucleus much the same as the planets aren’t pulled into the sun or the moon into the earth • According to classical mechanics (physics) charged particles moving in a curved path should emit energy (light) or some other form of electromagnetic radiation. Eventually, they would lose enough energy to be pulled into the nucleus Neutron In 1928, a German physicist, Walter Bothe, and his student, Herbert Becker, took the initial step in the search for the neutron. They bombarded beryllium with alpha particles emitted from polonium and found that it gave off a penetrating, electrically neutral radiation, which they interpreted to be high-energy gamma photons. Neutron cont’d • Eventually, the neutron was discovered in 1932 when James Chadwick used scattering data to calculate the mass of this neutral particle. – Chadwick is credited with his discovery Chadwick’s Experiment Basic Structure of Modern Atom There are three major subatomic particles: The proton with a stadardized charge of +1, a mass 1.673 x 10-24 g, and located in the nucleus The neutron with a charge of 0, a mass of 1.678 x 10-24 g, and located in the nucleus The neutron is slightly larger than the proton By what? Major Subatomic Particles cont’d The electron with a standardized charge of –1, a mass of 9.11 x 10-28 g, and located outside the nucleus in what is called the electron cloud The electron is ~ 1 / 1837 the mass of the proton Facts about the Basic Structure of Modern Atom Atoms are usually found in their neutral state, having equal numbers of protons and electrons # protons = # electrons for NEUTRAL atoms Most of the atom is composed of empty space (vacuum), the volume of the atom being established by the electron cloud Facts about the Basic Structure of Modern Atom The nucleus is an extremely small, extremely dense, positive core of the atom which contains the vast majority of the atom’s mass Nuclear density is about 1013 to 1014 g/cm3 Why don’t the electrons follow Newton’s First law and fly off and out of the atom? Electrostatic force - Coulombic attraction Facts about the Basic Structure of Modern Atom Atomic diameters are expressed on the order of angstroms 1 angstrom = 10-10 m Symbol for angstrom is Å Atomic nuclei have diameters that are on the order of 10-4 angstroms The identity of an element is established by the number of protons contained in the nuclei of its atoms This number of protons is referred to as the atomic number of the element The symbol for the atomic number is Z Facts about the Basic Structure of Modern Atom Atoms may lose or gain electrons without any change in element identity, BUT any change in number of protons will create a new identity for that atom – Loss of electrons results in ions with a positive charge (called cations) – Gain of electrons results in ions of a negative charge (called anions) Isotopes Isotopes are atoms of the same element which differ in mass This mass difference is due to possession of differing numbers of neutrons Recognize, nuclear charge cannot change Isotopes of an element have different mass numbers Mass number (symbol A) is equal to the sum of major nucleons A nucleon refers to any particle in the nucleus of an atom Mass Number Therefore, mass number equals the sum of protons and neutrons Mass number = Z + #n or A = Z + #n since A is the symbol for the mass number – To find the number of neutrons in an isotope use; #n = A - Z Representing Isotopes Isotopes have the general form of; AZX, where X is the chemical symbol of the element e.g. 1H; 14C; 206Pb often the value of Z isn’t expressed – why not? If the isotope is an ion a charge with a + or – sign would be in the right superscript e.g. 5726Fe+2 Naming Isotopes The term nuclide is used to refer to a specific atom of an isotope – an atom with specific number of protons and neutrons Hydrogen has three isotopes with special names: 1H - protium; 2H - deuterium; 3H - tritium Other naming involves the element name, a dash, and the mass number – e.g. 57Fe would be iron – 57 (implies neutral) – e.g. 57Fe+2 would be the iron – 57 +2 cation (or ion) – E.g. 33S-2 would be the sulfur – 33 -2 anion (or ion) Atomic Mass Scale Atoms are too small too measure their mass directly, however, their relative masses may be determined. Compounds are evaluated for the combining ratios of the atoms. After which upon decomposition, an analysis of the mass of each element may be established. From this information, the relative mass of each element may be determined. Atomic Mass Scale In order to provide uniformity, the atomic mass scale was established. The current reference standard is the isotope carbon – 12 (12C). It is assigned a mass of exactly 12 atomic mass units (a.m.u. or u.). 1 nuclide 12C = 12.000 000 000… a.m.u. What is an Atomic Mass Unit? The atomic mass unit is the unit of the atomic mass scale. It is defined as 1/12th the mass of the carbon – 12 nuclide. It is abbreviated as a.m.u. or u. 1 a.m.u. = 1/12 12C nuclide Average Atomic Mass In looking at the masses of the elements on the periodic table it is evident that most values listed aren’t close to being whole numbers Why? These values are actually averages that represent all the naturally occurring isotopes and the relative abundance in which they are found in nature Average Atomic Mass The average atomic mass may be determined by calculation from any analysis yielding the actual masses of the isotopes and some relative proportion of their presence (an actual number of nuclides or as percentages of nuclides) A “rough estimate” may be obtained if the mass numbers are utilized Average Atomic Mass Equations Ave. At. Mass = [(% x isotope mass) + (% x isotope mass) + …..] Total % Ave. At. Mass = [(#nuclides x isotope mass) + (#nuclides x isotope mass) + ..] Total # nuclides Mass Spec The mass spectrometer or mass spectrograph is a useful instrument in determining chemical analysis or analysis of isotopes The atoms of material are ionized giving them a positive charge Their charge is usually +1 but may sometimes be greater The ions travel through the magnetic field assuming a curved path which directs them onto a detector or film plate Number of hits at a particular spot gives the abundance (concentration) **Placement of hits leads to the calculation of mass to charge and eventually identity Mass Spec diagram Deflection of particles Mass Spec Different ions are deflected by the magnetic field by different amounts. The amount of deflection depends on: – the mass of the ion. Lighter ions are deflected more than heavier ones. – the charge on the ion. Ions with 2 (or more) positive charges are deflected more than ones with only 1 positive charge. These two factors are combined into the mass/charge ratio. Mass/charge ratio is given the symbol m/z (or sometimes m/e). – For example, if an ion had a mass of 28 and a charge of 1+, its mass/charge ratio would be 28. An ion with a mass of 56 and a charge of 2+ would also have a mass/charge ratio of 28. Mass Spec In the diagram above, ion stream A is most deflected - it will contain ions with the smallest mass/charge ratio. Ion stream C is the least deflected - it contains ions with the greatest mass/charge ratio. It makes it simpler to talk about this if we assume that the charge on all the ions is 1+. Most of the ions passing through the mass spectrometer will have a charge of 1+, so that the mass/charge ratio will be the same as the mass of the ion Note: You must be aware of the possibility of 2+ (etc) ions, but the vast majority of first level questions will give you mass spectra which only involve 1+ ions. Unless there is some hint in the question, you can reasonably assume that the ions you are talking about will have a charge of 1+. Mass Spec Assuming 1+ ions, stream A has the lightest ions, stream B the next lightest and stream C the heaviest. Lighter ions are going to be more deflected than heavy ones.