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UNIT 5 Periodic Properties Of The Elements

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UNIT 5
THE PERIODIC PROPERTIES OF THE ELEMENTS
Scientists constantly seek ways to organize facts so that they can identify similarities, differences, and trends among these facts. The most significant tool for organizing and remembering chemical facts is the Periodic Table.

5.1.

PERIODIC LAW: HISTORY and DEVELOPMENT of the PERIODIC TABLE

The discovery of new chemical elements has been an ongoing process since ancient times. Certain elements, such as gold and silver, appear in nature in elemental form and were thus discovered thousands of years ago. In contrast, some elements are radioactive and intrinsically unstable while majority of the elements, although stable, are dispersed widely in nature and are incorporated into numerous compounds. Thus for centuries, scientists were unaware of their existence. Finally, in the early th

19 century, advances in chemistry made it easier to isolate elements from their compounds. As a result, the number of elements known had more than doubled from 31 elements in 1800 to 63 elements in 1865.

As the number of known elements increased, scientists began to investigate the possibilities of classifying th
them in useful ways. Early in the 19 century, chemists became interested in the chemical and physical similarities between the elements.
TABLE 5-1: PRECURSORS OF THE PERIODIC TABLE
NAME OF
SCIENTISTS
Sir Humphry
Davy; and
Michael Faraday

Johann Wolfgang
Dobereiner

Beguyer de
Chancourtois
Robert Wilhelm
Bunsen; and
Gustav Robert
Kirchhoff
Stanislao
Cannizzaro

John A.R.
Newlands

DISCOVERIES
th
“Electrochemistry”. In the first quarter of the 19 century, scientists could determine the relative weights of atoms of the then known elements. The development of electrochemistry during this period by the British chemists Humphry Davy and Michael Faraday led to the discovery of many additional elements (Barium, Calcium, Boron, Potassium, Sodium, and Strontium).

“Dobereiner’s Triads”. As early as 1816, Dobereiner, a German chemist, noticed that there were similarities in the properties of several sets of three elements which he called “triads” (Li, Na, K; Ca, Sr, Ba; S, Se, Te; and Cl, Br, I). In each case, the atomic weight of the second or middle element in each set seemed to be approximately equal to the average of the atomic weights of the other two elements. He published this in an article in 1929.

“Telluric Helix”. In 1862, de Chancourtois, a French geologist, arranged the elements in order of atomic mass. When he wound the list spirally around a cylinder (like a modern helix), he found that similar elements fell along the same vertical lines.

“Spectroscopy”. The development of the spectroscope in 1859 by the German physicists Robert Wilhelm Bunsen and Gustav Robert Kirchhoff made possible the discovery of many more elements (Cesium, Ce and Rubidium, Rb). “Diatomic Elements”. In 1860, at the first international chemical congress ever held, the Italian chemist Stanislao Cannizzaro clarified the fact that some of the elements—for example, oxygen—have molecules containing two atoms. This realization finally enabled chemists to achieve a self-consistent listing of the elements.

“Law of Octaves”. In 1864 the British chemist John A. R. Newlands listed the elements in the order of increasing atomic weights and noted that a given set of properties recurs at every eighth place. He named this periodic repetition the law of octaves, by analogy with the musical scales. Newlands's discovery failed to impress his contemporaries, probably because the observed periodicity was limited to only a small number of the known elements.

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“Periodic Table”. The chemical law that the properties of all the elements are periodic functions of their atomic weights was developed independently by two chemists, in 1869 by the Russian chemist Dmitri Mendeleyev and in 1870 by the German chemist Julius Lothar Meyer.

Dmitri Ivanovich
Mendeleev

In 1869, Dmitri Mendeleev proposed a Periodic Law – “when the elements are arranged in order of increasing atomic mass, similarities in properties recur periodically”.
Mendeleev also developed the first Periodic Table of Elements in 1869. He listed the elements in such a way that similar elements appeared in vertical columns. In order to make similar elements appear under one another. Mendeleev had to leave blanks for undiscovered elements in his table. The subsequent discovery of Scandium, Gallium and Germanium, each of which was to have properties much like those predicted by Mendeleev, demonstrated the validity of the periodic system.

The existence of the Nobel Gases (Inert Gases) was unforeseen by Mendeleev, nevertheless, after their discovery in the years 1894 – 1898, these elements readily fitted into the periodic table.

Julius Lothar
Meyer

John William
Strutt, 3rd Baron
Rayleigh;
and
Sir William
Ramsay.

Henry
G.
Moseley

J.

Mendeleev also had to place certain elements (K, Ni, and I) out of order to get them into the proper groups of his periodic table. He assumed this was because of errors in atomic masses. It was not until 1913 that theoretical explanation for this reordering was discovered.

“Periodic Table”. In 1868 Lothar Meyer, a German chemist, came up independently with an arrangement of elements similar to that of Mendeleev. Many believe that he should share the credit for the periodic table. Unfortunately, Meyer did not write up his table until December 1869, and it was not published until March 1870 while Mendeleev had published his table in 1869.

“Noble Gases”. The periodic law has undergone two principal elaborations since its original formulation by Mendeleyev and Meyer. The first revision involved extending the law to include a whole new family of elements, the existence of which was completely unsuspected in the 19th century. This group comprised the first three of the Noble, or Inert, Gases: Argon, Helium, and Neon, discovered in the atmosphere between 1894 and 1898 by the British physicist John William Strutt and the British chemist Sir William Ramsay.

The second development in the periodic law was the interpretation of the cause of the periodicity of the elements in terms of the Bohr theory (1913) of the electronic structure of the atom by Henry Moseley.

“Atomic Numbers”. Between 1913 and 1914, as a result of his research on the Xray spectra of 38 elements with atomic numbers between 13 (Aluminum, Al) and 79 (Gold, Au), H.J.G. Moseley, a British physicist, was able to solve the problem about the relationship of the fundamental properties of the periodic table and atomic number.

Using a corresponding spectral line for each element, he found that there is a linear relationship between the square root of the frequency of the line and the atomic number of the element. In other words, the square root of the frequency of the spectral line increases by a constant amount from element to element when the elements are arranged by increasing atomic number.

Moseley was able to assign correct atomic number to any element on the basis of its X-ray spectrum. He settled the problem involving the classification of elements that have atomic weights out of line with those of their neighbors (Potassium, Nickel, and Iodine).

He also stated that there should be 14 elements in the series from element 58, Cerium to element 71, Lutetium. He also established that these elements should follow Lanthanum in the periodic table.

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Modern Periodic Law – states that “the chemical and physical properties of the elements are periodic functions of Atomic Number”.
Periodic Table – is a very useful device for correlating the fundamental properties of the different elements. It is designed in such a way that similar elements are grouped together and the properties of the elements can be predicted from their positions in the table. 5.2.

THE CLASSIFICATION OF THE ELEMENTS
According to their Electron Configurations:

5.2.1. The Noble Gases – are the colorless, monatomic gases which are chemically un-reactive and diamagnetic. These are the colorless gases that have similar chemical and physical properties and of low reactivity. They possess the most stable electron configurations. They have atomic numbers 2, 10, 18, 36, 54, and 86 respectively. In the periodic table, the noble gases are found at the end of each period in Group VIII-A or Group O ( Helium, Neon, Argon, Krypton, Xenon, and Radon).

5.2.2. The Representative Elements – these elements are found in the “A – Families” of the periodic table and include both metals and nonmetals. They exhibit a wide range of chemical and physical characteristics. Some of the elements are diamagnetic and some are paramagnetic, but the compounds of these elements however, are generally diamagnetic and colorless. All their electron shells are either complete or stable except the outer shell to which the last electron may be considered as having been added. This outer shell is termed the valence shell and electrons in it are the valence electrons. The number of the valence electrons for each atom is the same as the “group number”. The chemistry of these elements depends upon these valence electrons.

Diamagnetic Elements – are elements that are repelled by a magnetic field. In such substances, all their electrons are paired.
Paramagnetic Elements – are the elements that are drawn into a magnetic field. Such substances contain unpaired electrons.

5.2.3. The Transition Elements – these elements are found in the “B – Families” of the periodic table. 5.2.4. The Inner Transition Elements – These elements are found at the bottom of the periodic table, th

th

but they belong to the 6 and 7 periods after the elements in Group III-B. All inner transition elements are metals. They are paramagnetic and their compounds are paramagnetic and colored.
The Rare Earth Elements – (Rare Earth Metals) are the series of elements with atomic numbers 57 through 71. They are named in order: Lanthanum, Cerium, Praseodymium, Neodymium, Promethium, Samarium, Europium, Gadolinium, Terbium, Dysprosium, Holmium, Erbium, Tulium, Yterbium, and Lutetium. Yttrium (atomic no. 39) and scandium (atomic no. 21) are sometimes included in the group of rare earth elements. The elements cerium (atomic no. 58) through lutetium (atomic no. 71) are also commonly known as the “Lanthanide Series of Elements”, named after the element “Lanthanum” which precedes them.

The Actinide Series of Elements – (The Actinoids) a series of 14 radioactive elements in the periodic table with atomic numbers 89 through 102. Only the first four elements in the series have been found in nature in appreciable amounts; the remainders have been produced synthetically. Those elements with atomic numbers of 93 and above are called transuranium elements. The elements constituting the actinide series are, in order of increasing atomic number: Actinium, Thorium, Protactinium, Uranium, Neptunium, Plutonium, Americium, Curium, Berkelium, Californium, Einsteinium, Fermium, Mendelevium, Nobelium, and Lawrencium.

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The Transuranium Elements – These are the chemical elements with atomic number greater than 92, the atomic number of uranium in the periodic table. More than 20 such elements have been identified. These elements consist of more than 100 radioactive isotopes, which are characterized by radioactive instability. These radioisotopes are produced artificially by bombarding heavy atoms either with neutrons produced in nuclear reactors or with charged particles accelerated to high energy in particle accelerators. The first 10 transuranium elements, together with actinium, thorium, protactinium, and uranium, constitute the “actinide elements”, which are chemically analogous to the rare earth elements. They are, in order: Neptunium, Plutonium, Americium, Curium, Berkelium, Californium, Einsteinium, Fermium, Mendelevium, Nobelium, and Lawrencium.

5.3.

FEATURES OF THE PERIODIC TABLE

5.3.1. PERIOD – is the collective name of all the elements found and arranged in a “Horizontal Row” in the table.
st

The 1 Period consists of “Only Two (2) Elements: Hydrogen (H) and Helium (He). nd

rd

nd

The 2
and 3 Periods respectively are consists of Eight (8) Elements each. The 2 period rd
runs from Lithium (Li) to Neon (Ne), while the 3 period starts with Sodium (Na) and ends with Argon (Ar).
The subsequent Periods contain 18, 18, 32 and 32 elements respectively. st

With the exception of the 1 period, each period starts with an Alkali Metal and ends with a Noble Gas.

5.3.2. GROUP or FAMILY – is the collective name of the elements that appear in a “Vertical Column” in the table.
The elements of the group have similar chemical properties. Example are the Nobel Gases, Alkali Metals, Alkaline Earth Metals and the Halogens.
Each group is given a designation that usually consists of a “Roman Numeral” followed by the letter “A” or “B”.

5.3.3. GROUP TRENDS
A. THE ACTIVE METALS:
1. Alkali Metals – are the group of elements that consists of highly reactive, soft-metallic solids. All these group I-A elements have characteristics metallic properties such as a silvery, metallic luster and high thermal and electrical conductivities. The name “Alkali” comes from the Arabic word meaning “ashes”. These are composed of the elements: Lithium (Li, Z = 3), Sodium (Na, Z = 11), Potassium (K, Z = 19), Rubidium (Rb, Z = 37), Cesium (Ce, Z = 55) , and Francium (Fr, Z = 87). The two most abundant alkali metals , were isolated from wood ashes by the early chemists. The names “Soda Ash” and “Potash” are still sometimes used for the carbonate salts Sodium Carbonate (Na2CO3) and Potassium Carbonate (K2CO3).

2. Alkaline Metals – are another group of elements similar to the alkali metals. These group II-A elements are all solids with typical metallic properties. Compared with the alkali metals, these metals are harder, denser, and melt at higher temperatures. The Alkaline Earth Metals namely: Beryllium (Be, Z = 4), Magnesium (Mg, Z = 12), Calcium (Ca, Z = 20), Strontium (Sr, Z = 38) and Barium (Ba, Z = 56) are less reactive than the alkali metals.

4

B. SOME SELECTED NONMETALS:
1

1. Hydrogen – the first element in the periodic table has a 1s electron configuration and is usually placed above the alkali metals. However, it is a unique element and does not truly belong to any family or group. Hydrogen, the lightest element on earth, is a non-metal that occurs as a colorless diatomic gas, H2 (Hydrogen gas). Hydrogen generally reacts with other non-metals to form molecular compounds and it is also reactive with active metals to form solid metal hydrides, e.g. NaH.

2. Oxygen Family – are the group VI-A elements and are mostly nonmetals (Oxygen, Sulfur and Selenium) with the exception of Tellurium – a metalloid, and Polonium – a metal, which is radioactive and quite rare. The first element in the group, Oxygen, is a colorless gas at room temperature and is encountered in two molecular forms, O2 – Oxygen gas and O3 – Ozone. The two forms of oxygen are called “Allotropes – are different forms of the same element in the same state”.

3. Halogens – are the elements before the Noble gases (group VIII-A) in each complete st
period (except for Helium in the 1 period), and they are very reactive nonmetals. The Halogens are the elements under group VII-A composed of: Fluorine (F), Chlorine (Cl), Bromine (Br), and Iodine (I). The group 7A elements are known as Halogens, after the Greek words “halos” and “gennao”, meaning “Salt Formers”. All the halogens are typical nonmetals. Their melting and boiling points increase with increasing atomic number. Fluorine and Chlorine are gases at room temperature, Bromine is a liquid, and Iodine is a solid. Each element consists of diatomic molecules: F2, Cl2, Br2, and I2. Fluorine gas is pale yellow, Chlorine gas has a yellow-green color; Bromine liquid is reddish brown and readily forms a reddish-brown vapor; and solid Iodine is grayish black and readily sublimes into a violet vapor.

4. Noble Gases – are the group VIII-A elements and are all nonmetals that are gases at room temperature. The noble gases are characterized by completely filled s and p- orbitals. All elements of group VIII-A have large first ionization energies, and because the noble gases posses such stable electronic configurations, they are exceptionally un-reactive. In fact, until the early 1960’s the elements were called “Inert Gases”, until Neil Bartlett from the University of British Columbia, synthesized the first Noble gas compound by reacting Xenon with Fluorine containing compound Platinum Hexafluoride (PtF6), to form the compounds Xenon Difluoride, XeF2; Xenon Tetrafluoride, XeF4, and Xenon Hexafluoride, XeF6.

5.3.4. METALS, NON-METALS and METALLOIDS:
A. Metals – are the elements that in general has a characteristic luster, conducts heat and electricity well, and can be pounded and stretched into shapes without breaking (malleable and ductile). About 80% of the known elements are metals.

B. Nonmetals – are the elements that are not lustrous, are poor conductor or electricity and heat, and are brittle in the solid state. Its chemical properties differ from those of the metals. They are placed in the left-hand side of the stepped-diagonal line found in the periodic table. NOTE: The “Stepped Diagonal Line” found in the Periodic Table marks the approximate division between the metals (found in the Right-hand side of this line) and the non-metals (found in the Left-hand side of the line).

C. Metalloids – are the elements that fall close on both sides of the line division separating nonmetals from metals and therefore have properties that are intermediate between those of the metals and nonmetals. These are the elements: Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Selenium (Se), Antimony, and Tellurium. The metalloids or “semiconducting elements” are mostly brittle solids with electrical conductivity intermediate between that of metals, which are good conductors, and nonmetals, which are nonconductors. These electrical properties of metalloids make them essential components for computer technology and electronic industry.

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5.4.

ATOMIC RADII, IONIZATION ENERGY and ELECTRON AFFINITY

The size of an atom is not a definite quantity; however, the boundary is somewhat fuzzy and may vary. For these reasons and also because it is not possible to measure the size of an isolated atom , chemists compare the sizes by measuring the distance between atoms in an element. A. The ATOMIC RADIUS – is one-half the distance between the nuclei of identical atoms joined in a molecule. For example, the distance between the nuclei of two Iron (Fe) atoms (248 pm) in a sample of solid iron can be measured. The radius of an iron atom is then taken as one-half of this distance (124 pm).

ATOMIC RADIUS VARIATION in the PERIODIC TABLE:
The size of the atoms of an element varies in a regular way across the periodic table, increasing down the groups (columns), and decreasing along the periods (rows) from left to right. Group Trends:

The size of an atom is largely determined by its electrons. The electrons are arranged in shells surrounding the nucleus of each atom. The top elements of every group have only one or two electron shells. Atoms of elements further down the table have more shells and are therefore larger in size.

Period Trends:
The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus (number of protons). Moving across a period from left to right, the outermost electron shell fills up but no new shells are added. At the same time, the number of protons in the nucleus of each atom increases. Protons attract electrons. The greater the number of protons present, the stronger the attraction that holds the electrons closer to the nucleus, and the smaller the size of the shells.

B. ION – is an atom or group of atoms that has a positive or negative charge. Sodium, for example, +1
readily forms an Na ion. Any process that results in the formation of an ion is referred to as “Ionization”.
Ionization Energy – is the energy required to remove one electron from an atom of an element.

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C. ELLECTRON AFFINITY – is the energy change that occurs when an electron is acquired by a neutral atom. Neutral atoms can acquire electrons. The energy involved when an atom gains an electron is referred to as the electron affinity.

Many atoms readily add electrons and release energy, which classifies this as an exothermic process and by convention, the quantity of energy released in an exothermic process is represented by a “negative number”.

Some atoms, however, must be ―forced‖ to gain an electron by the addition of energy, and the quantity of energy absorbed in such an endothermic process is represented by a “positive number”. An ion produced in this way is likely to be unstable and to lose the added electron spontaneously.

5.5.

THE NAMES AND SYMBOLS OF THE ELEMENTS

Chemists talk and write about elements so frequently that they have introduced a set of symbols to represent the elements. These constitute a firm of chemical shorthand. The symbols used fall into four categories:

TABLE 5-2: THE NAMES AND SYMBOLS OF THE ELEMENTS
Examples of the Elements
Categories:

1.) For some elements we use a
symbol derived from their
original “Latin name”.

2.) Some elements are
symbolized by simply using
the “first letter” of their
names (always a capital
letter). These symbols are
generally used for very
common element.

3.) Where there are several
elements beginning with the
same letter, a “second
letter” is used. The first
method is to symbolize the
elements using the “first two
letters” of their name. The
first letter is written as a
capital letter (upper case)
and the second letter a
small letter (lower case).

Modern Name
Sodium
Potassium
Iron
Copper
Silver
Tin
Antimony
Tungsten
Gold
Mercury
Lead
Hydrogen
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Phosphorus
Sulfur

Natrium
Kalium
Ferrum
Cuprum
Argentum
Stannum
Stibium
Wolfram
Aurum
Hydrargyrum
Plumbum
Hydro (water) + genes
Borax
“Carbo” – Charcoal
Nitron (Nitre) + genes
Oxys + genes
Fluo – “flow”
“Light Bearing”
Sulphur– Brimstone

Chemical
Symbol
Na
K
Fe
Cu
Ag
Sn
Sb
W
Au
Hg
Pb
H
B
C
N
O
F
P
S

Lithium
Aluminum
Silicon
Calcium
Titanium
Cobalt
Nickel
Germanium
Bromine
Indium
Xenon
Barium

“Lithos” – or Stone
Alumen or Alum
“Silix” – or Flint
“Calx” – or Lime
Titan – Greek Mythology
Kobold – Evil spirit
Kupfernickel – false nickel
Named after Germany
“Bromos” – or Stench
Indigo – color
“Xenos” – Stranger
“Barys” – Heavy or Dense

Li
Al
Si
Ca
Ti
Co
Ni
Ge
Br
In
Xe
Ba

Origin of Name

7

4.) Other elements are
symbolized using the “first
two consonants” or “first
syllable in succession” of
their name. The first letter is
again written as a “capital
letter” followed by the
second letter in “small
letter”.

5.6.

Magnesium
Chlorine
Chromium
Manganese
Zinc
Arsenic
Rubidium
Zirconium
Platinum
Astatine
Plutonium

Magnesia – Ancient City
“Chloros” – greenish yellow
“Chroma” – color in Latin
“Magnes” – Magnet
“Zin” – German for Tin
“Arsenikos” – Male
“Rubidius” – or Red
“Zicon” – a Gem
“Platina” – Little silver
“Astatos” – Unstable
Pluto – The Planet

Mg
Cl
Cr
Mn
Zn
As
Rb
Zr
Pt
At
Pu

TYPES OF BONDING AND THE VALENCY OF THE ELEMENTS

The type of bonding associated with elements in some of the groups of the Periodic Table: 1. The Alkali Metals (Group I-A) all tend to lose one (1) electron and therefore form “Ionic +1
+1
+1
+1
+1
Compounds” involving the single charged ions, Li , Na , K , Rb , and Cs . 2. The Alkaline Earth Metals (Group II-A) all tend to lose two (2) electrons and therefore form +2
+2
+2
+2
+2
ionic compounds, this time with double charged ions: Be , Mg , Ca , Sr , and Ba . 3. The Halogen Nonmetals (Group VII-A) all tend to gain one (1) electron. They can do this by forming either ionic compounds (when they combine with the alkali or alkaline elements which want to lose electrons) or “covalent compounds (when they combine with other nonmetals which also want to gain electrons)”

4. The Oxygen Family (Group VI-A) similarly want to gain electrons and thus can form both ionic (as in K2O and CaO) and covalent compounds (as in H2O and SO2). 5. The Nitrogen and Carbon Family (Groups IV-A and V-A) almost always form covalent compounds (as in NH3, PCl3, CH4, and SiF4).

6. In Group III-A, Boron usually forms covalent compounds, while Aluminum forms ionic +3
compounds (involving the Al ion).
+3

+3

+2

7. The Transition Metals all lose electrons to form positively charged ions (e.g. Cr , Fe , Cu , +1
+2
Ag , Zn ). But it is not possible from a simple look at the Periodic table to predict just how many electrons any particular atom will lose. In addition, the transition metals can form covalent bonds with Oxygen to form what are called “Oxyacids” (e.g. H2NO3, H2SO4, H3PO4, H2CO3) and “Oxyanions” (e.g. ClO2, NO3, ZnO2, SO4, BrO3, PO4, CO3). The Valency or Valence of an Element – is a number which serves as a measure of the combining capacity of the element when it forms compounds.

When an element forms “Ionic Compounds”, the valence of the element is the charge the atom (ion) carries.
When an element forms “Covalent Compounds”, the valence of the element is the number of covalent bonds that the atom forms (single bond, counted as one; double bonds, counted as two; and triple bonds counted as three).

_____________________________________________________________________________ Pb: Ms. Miroma R. Villacrucis - Instructor / UNIT 5 - The Periodic Properties of the Elements / July 12, 2013

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